Chemistry I Notes Unit 3 Chapters 7-9

Chemistry I Notes Unit 3 Chapters 7-9
Chemical Formulas, Bonding, and Naming Big Idea #3 Bonding and Interactions

Valence Electrons Valence electrons are the electrons in the highest occupied energy level of an element’s atoms. The number of valence electrons largely determines the chemical properties of an element. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Valence Electrons Determining the Number of Valence Electrons
To find the number of valence electrons in an atom of a representative element, simply look at its group number. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Atoms of the Group 1A elements (hydrogen, lithium, sodium, and so forth) all have one valence electron, corresponding to the 1 in 1A. Carbon and silicon atoms, in Group 4A, have four valence electrons. The noble gases (Group 8A) are the only exceptions to the group-number rule: Atoms of helium have two valence electrons, and atoms of all the other noble gases have eight valence electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

7.1 Ions – Atoms that have lost or gained one or more electrons.
Positive ions have lost electrons (e-) Negative ions have gained electrons (e-) The symbol indicates the number of electrons lost or gained. How many protons neutrons and electrons do these contain?

Lewis Dot Structures & Ions

Valence Electrons The Octet Rule
Noble gases, such as neon and argon, are nonreactive in chemical reactions. That is, they are stable. In 1916, chemist Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules. He called his explanation the octet rule. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The octet rule states that in forming compounds, atoms tend to achieve the electron configuration of a noble gas. An octet is a set of eight. Atoms of each of the noble gases (except helium) have eight electrons in their highest occupied energy levels and the general electron configuration of ns2np6. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Atoms of metals tend to lose their valence electrons, leaving a complete octet in the next-lowest energy level. Atoms of some nonmetals tend to gain electrons or share electrons with another nonmetal atom or atoms to achieve a complete octet. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Draw the electron dot structure for bismuth.
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Formation of Cations A positively charged ion, or cation, is produced when an atom loses one or more valence electrons. A sodium atom (Na) forms a sodium cation (Na+). A calcium atom (Ca) forms a calcium cation (Ca2+). Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Cations Group 1A Cations
The most common cations are those produced by the loss of valence electrons from metal atoms. Most of these atoms have one to three valence electrons, which are easily removed. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

When forming a compound, a sodium atom loses its one valence electron and is left with an octet in what is now its highest occupied energy level. The number of protons in the sodium nucleus is still eleven, so the loss of one unit of negative charge produces a cation with a charge of 1+. Na 1s22s22p63s Na+ 1s22s22p6 –e– octet Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Magnesium atom (electrically neutral, charge = 0)
Formation of Cations Group 2A Cations Magnesium (atomic number 12) belongs to Group 2A of the periodic table, so magnesium atoms have two valence electrons. • Magnesium atom (electrically neutral, charge = 0) • Mg Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Magnesium (atomic number 12) belongs to Group 2A of the periodic table, so magnesium atoms have two valence electrons. A magnesium atom attains the electron configuration of a neon atom by losing both valence electrons and producing a magnesium cation with a charge of 2+. • Mg Mg e– • loses all its valence electrons Magnesium atom (electrically neutral, charge = 0) Magnesium ion (2+ indicates two units of positive charge) (2 in front of e– indicates two units of negative charge) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Cations The figure at right lists the symbols of the cations formed by metals in Groups 1A and 2A. Cations of Group 1A elements always have a charge of 1+. Cations of Group 2A elements always have a charge of 2+. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Cations Transition Metal Cations
The charges of cations of the transition metals may vary. An atom of iron may lose two valence electrons, forming the Fe2+ cation, or three valence electrons, forming the Fe3+ cation. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Some ions formed by transition metals do not have noble-gas electron configurations (ns2np6) and are therefore exceptions to the octet rule. Silver, with the electron configuration of 1s22s22p63s22p63d104s24p64d105s1, is an example. To achieve the structure of krypton, a silver atom would have to lose eleven electrons. To acquire the electron configuration of xenon, a silver atom would have to gain seven electrons. Ions with charges of three or greater are uncommon. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

A copper atom loses its lone 4s electron to form a copper ion (Cu+) with a pseudo noble-gas electron configuration, as illustrated below. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Anions Formation of Anions How are anions formed?

Formation of Anions An anion is produced when an atom gains one or more valence electrons. Note that the name of an anion of a nonmetallic element is not the same as the element name. The name of the anion typically ends in -ide. Thus, a chlorine atom (Cl) forms a chloride anion (Cl–). An oxygen atom (O) forms an oxide anion (O2–). Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Anions Atoms of nonmetals and metalloids form anions by gaining enough valence electrons to attain the electron configuration of the nearest noble gas. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Anions Atoms of nonmetallic elements attain noble-gas electron configurations more easily by gaining electrons than by losing them because these atoms have relatively full valence shells. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Anions Atoms of nonmetallic elements attain noble-gas electron configurations more easily by gaining electrons than by losing them because these atoms have relatively full valence shells. Atoms of chlorine have seven valence electrons. A gain of one electron gives a chlorine atom an octet and converts a chlorine atom into a chloride atom. Cl 1s22s22p63s23p Cl– 1s22s22p63s23p6 +e– octet Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Anions Chlorine atoms need one more valence electron to achieve the electron configuration of the nearest noble gas. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Anions The ions produced when atoms of chlorine and other halogens gain electrons are called halide ions. All halogen atoms have seven valence electrons and need to gain only one electron to achieve the electron configuration of a noble gas. All halide ions (F–, Cl–, Br–, and I–) have a charge of 1–. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formation of Anions Oxygen is in Group 6A, and an oxygen atom has six valence electrons. An oxygen atom attains the electron configuration of neon by gaining two electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

This table lists some common anions.
Interpret Data This table lists some common anions. Some Common Anions Name Symbol Charge Fluoride F– 1– Chloride Cl– Bromide Br– Iodide I– Oxide O2– 2– Sulfide S2– Nitride N3– 3– Phosphide P3– Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

What is the electron configuration of a sulfide ion
What is the electron configuration of a sulfide ion? What noble gas shares this configuration? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

7-2 Ionic Bonding It is possible that, as two atoms come close, one electron is transferred to the other atom. The atom that gives up an electron acquires a +1 charge and the other atom, which accepts the electron acquires a –1 charge. The two atoms are attracted to each other through Coulombic interactions – opposite charges attract – resulting in an IONIC bond. Animation

Formation of Ionic Compounds
Ionic Bonds Anions and cations have opposite charges and attract one another by means of electrostatic forces. The electrostatic forces that hold ions together in ionic compounds are called ionic bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Ionic Bonds When sodium and chlorine react to form a compound, the sodium atom transfers its one valence electron to the chlorine atom. Sodium and chlorine atoms combine in a one-to-one ratio, and both ions have stable octets. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Ionic Bonds Aluminum metal (Al) and the nonmetal bromine (Br2) react violently to form the ionic solid aluminum bromide (AlBr3). Each bromine atom has seven valence electrons and readily gains one additional electron. Three bromine atoms combine with each aluminum atom. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formula Units A chemical formula shows the numbers of atoms of each element in the smallest representative unit of a substance. NaCl is the chemical formula for sodium chloride. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Generally, Metallic element + nonmetallic element  IONIC compound Ionic compounds are generally high-melting point solids that are good conductors of heat and electricity in the molten state. Examples are NaCl, common salt, and NaF, sodium fluoride.

Formula Units Ionic compounds do not exist as discrete units, but as collections of positively and negatively charged ions arranged in repeating patterns. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formula Units The chemical formula of an ionic compound refers to a ratio known as a formula unit. A formula unit is the lowest whole-number ratio of ions in an ionic compound. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Formula Units For sodium chloride, the lowest whole-number ratio of the ions is 1:1 (one Na+ ion to each Cl– ion). The formula unit for sodium chloride is NaCl. Although ionic charges are used to derive the correct formula, they are not shown when you write the formula unit of the compound. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Predicting Formulas of Ionic Compounds
Sample Problem 7.1 Predicting Formulas of Ionic Compounds Use electron dot structures to predict the formulas of the ionic compounds formed from the following elements: a. potassium and oxygen b. magnesium and nitrogen Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Use electron dot structures to determine the formula of the ionic compound formed when calcium reacts with fluorine. • • F • + Ca2+ – Ca CaF2 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Properties of Ionic Compounds
Crystalline Solids at room temperature High melting points Electrolytes – conduct electricity in water solutions and when molten Coordination number is the number of oppositely charged ions surrounding a given ion in a crystal lattice.

What are the coordination numbers?

The coordination number of an ion is the number of ions of opposite charge that surround the ion in a crystal. In NaCl, each ion has a coordination number of 6. The coordination number of Na+ is 6 because each Na+ ion is surrounded by six Cl– ions. The coordination number of Cl– is also 6 because each Cl– ion is surrounded by six Na+ ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

In CsCl, each ion has a coordination number of 8. Each Cs+ ion is surrounded by eight Cl– ions. Each Cl– ion is surrounded by eight Cs+ ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Titanium dioxide (TiO2), or rutile, forms tetragonal crystals. The coordination number for the cation (Ti4+) is 6. Each Ti4+ ion is surrounded by six O2– ions. The coordination number of the anion (O2–) is 3. Each O2– ion is surrounded by three Ti4+ ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Most ionic compounds are crystalline solids at room temperature. The component ions in such crystals are arranged in repeating three-dimensional patterns. The beauty of crystalline solids comes from the orderly arrangement of their component ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

When sodium chloride is melted, the orderly crystal structure breaks down. To (+) electrode To (–) electrode Inert metal electrodes Flow of electrons Current meter Power source Cl– Na+ If a voltage is applied across this molten mass, cations migrate freely to one electrode and anions migrate to the other. This movement of electrons allows electric current to flow between the electrodes through an external wire. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Ionic compounds also conduct electric current if they are dissolved in water. When dissolved, the ions are free to move about in the solution. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

When can ionic compounds conduct an electric current?
A. Only when melted B. When melted or dissolved in water C. Only when dissolved in water D. When solid or melted Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

7.3 Bonding in Metals Valence electrons of atoms in a pure metal can be modeled as a sea of electrons surrounding the positively charged kernels. These “free” electrons can pass from one part of the metal to another. They surround and hold the positive cations together. Metallic bonds are the forces between the electron sea and the cations they surround. Metals can conduct electricity, are malleable (can be pounded into sheets), ductile (can be pulled into wires), and form a crystal lattice. Alloys are mixtures of 2 or more elements where at least one is a metal. Examples: Bronze, Brass, Pewter, and Steel

Metallic Bonds and Metallic Properties
The valence electrons of atoms in a pure metal can be modeled as a sea of electrons. The valence electrons are mobile and can drift freely from one part of the metal to another. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Metallic bonds are the forces of attraction between the free-floating valence electrons and the positively charged metal ions. These bonds hold metals together. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Properties of Metals Metals are good conductors of electric current because electrons can flow freely in the metal. As electrons enter one end of a bar of metal, an equal number of electrons leave the other end. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Properties of Metals Metals are ductile—that is, they can be drawn into wires. Force Metal rod Die Wire Metals are also malleable, which means that they can be hammered or pressed into shapes. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Properties of Metals When a metal is subjected to pressure, the metal cations easily slide past one another. Force Sea of electrons Metal cation Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Metal

Properties of Metals When a metal is subjected to pressure, the metal cations easily slide past one another. Sea of electrons Metal cation Force Strong repulsions Nonmetal anion Metal Ionic crystal If an ionic crystal is struck with a hammer, the blow tends to push the positive ions close together. The positive ions repel one another, and the crystal shatters. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Crystalline Structure of Metals For spheres of identical size, such as metal atoms, several closely packed arrangements are possible. These Thai oranges illustrate a pattern called a hexagonal close-packed arrangement. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Crystalline Structure of Metals In a body-centered cubic structure, every atom (except those on the surface) has eight neighbors. Chromium The metallic elements sodium, potassium, iron, chromium, and tungsten crystallize in a body-centered cubic pattern. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Crystalline Structure of Metals In a face-centered cubic arrangement, every atom has twelve neighbors. Among the metals that form a face-centered cubic structure are copper, silver, gold, aluminum, and lead. Gold Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Crystalline Structure of Metals In a hexagonal close-packed arrangement, every atom also has twelve neighbors. Zinc The pattern is different from the face-centered cubic arrangement. Metals that have a hexagonal close-packed crystal structure include magnesium, zinc, and cadmium. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Alloys Alloys Why are alloys important?
Alloys are mixtures of two or more elements, at least one of which is a metal. Brass, for example, is an alloy of copper and zinc. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Alloys Alloys are important because their properties are often superior to those of their component elements. Sterling silver (92.5 percent silver and 7.5 percent copper) is harder and more durable than pure silver, yet it is still soft enough to be made into jewelry and tableware. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Alloys The most important alloys today are steels.
The principal elements in most steels, in addition to iron and carbon, are boron, chromium, manganese, molybdenum, nickel, tungsten, and vanadium. Stainless Steel 80.6% Fe 18.0% Cr 0.4% C 1.0% Ni Steels have a wide range of useful properties, such as corrosion resistance, ductility, hardness, and toughness. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Alloys Alloys can form from their component atoms in different ways.

If the atoms of the components in an alloy are about the same size, they can replace each other in the crystal. This type of alloy is called a substitutional alloy. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

If the atoms of the components in an alloy are about the same size, they can replace each other in the crystal. This type of alloy is called a substitutional alloy. If the atomic sizes are quite different, the smaller atoms can fit into the interstices (spaces) between the larger atoms. Such an alloy is called an interstitial alloy. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

8.1 Covalent Bonding A chemical bond results from strong electrostatic interactions between two atoms. The nature of the atoms determines the kind of bond. COVALENT BONDS result from a strong interaction between NEUTRAL atoms Each atom donates an electron resulting in a pair of electrons that are SHARED between the two atoms Covalently bonded compounds are called Molecular Compounds because they form discrete molecules.

For example, consider a hydrogen molecule, H2
For example, consider a hydrogen molecule, H2. When the two hydrogen, H, atoms are far apart from each other they do not feel any interaction. As they come closer each “feels” the presence of the other. The electron on each H atom occupies a volume that covers both H atoms and a COVALENT BOND is formed. Once the bond has been formed, the two electrons are shared by BOTH H atoms. When a molecule is formed from two atoms of the same type it is called a diatomic molecule. Examples of elements that form diatomic molecules besides H2 are : O2, F2, Br2, I2, N2, Cl These element symbols spell HOFBrINCl

An electron density plot for the H2 molecule shows that the shared electrons occupy a volume equally distributed over BOTH H atoms. Electron Density for the H2 molecule

What factors determine if an atom forms a covalent or ionic bond with another atom?
The number of electrons in an atom, particularly the number of the electrons furthest away from the nucleus determines the atom’s reactivity and hence its tendency to form covalent or ionic bonds. These outermost electrons are the one’s that are more likely to “feel” the presence of other atoms and hence the one’s involved in bonding i.e. in reactions. Chemistry of an element depends almost entirely on the number of valence electrons, and hence its atomic number. Molecular Formulas show the number of atoms of each element in one molecule of a covalent compound. Empirical Formulas show the minimum ratio of the elements in a compound. This is the formula unit of an ionic compound

Molecules and Molecular Compounds
Ionic compounds are generally crystalline solids with high melting points. Other compounds, however, have very different properties. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Ionic compounds are generally crystalline solids with high melting points. Other compounds, however, have very different properties. Water (H2O) is a liquid at room temperature. Carbon dioxide (CO2) and nitrous oxide (N2O) are both gases at room temperature. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The attractions that hold together the atoms in O2, H2O, CO2, and N2O cannot be explained by ionic bonding. These bonds do not involve the transfer of electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sharing Electrons Recall that ionic bonds form when the combining atoms give up or accept electrons. Another way that atoms can combine is by sharing electrons. Atoms that are held together by sharing electrons are joined by a covalent bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sharing Electrons A molecule is a neutral group of atoms joined together by covalent bonds. Oxygen gas consists of oxygen molecules; each oxygen molecule consists of two covalently bonded oxygen atoms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sharing Electrons An oxygen molecule is an example of a diatomic molecule—a molecule that contains two atoms. Other elements found in nature in the form of diatomic molecules include hydrogen, nitrogen, and the halogens. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sharing Electrons Molecules can also be made of atoms of different elements. A compound composed of molecules is called a molecular compound. Water is an example of a molecular compound. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Representing Molecules The molecular formula of water is H2O. Notice that the subscript written after an element’s symbol indicates the number of atoms of each element in the molecule. If there is only one atom, the subscript 1 is omitted. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Representing Molecules Butane is also a molecular compound. Butane is commonly used in lighters and household torches. The molecular formula for butane is C4H10. According to this formula, one molecule of butane contains four carbon atoms and ten hydrogen atoms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Representing Molecules A molecular formula reflects the actual number of atoms in each molecule. The subscripts are not necessarily the lowest whole-number ratios. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Representing Molecules Note that molecular formulas also describe molecules consisting of atoms of one element. For example, an oxygen molecule consists of two oxygen atoms bonded together; its molecular formula is O2. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Representing Molecules A molecular formula does not tell you about a molecule’s structure. In other words, it does not show either the arrangement of the various atoms in space or which atoms are covalently bonded to one another. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Representing Molecules A variety of diagrams and molecular models can be used to show the arrangement of atoms in a molecule. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Representing Molecules The arrangement of atoms within a molecule is called its molecular structure. The molecular structure of carbon dioxide shows how the three atoms are arranged in a row. It also shows how the carbon atom in each molecule is in the middle between the two oxygen atoms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Representing Molecules The arrangement of atoms within a molecule is called its molecular structure. The molecular structure of water shows how the oxygen atom is in the middle between the hydrogen atoms. The atoms in water are not arranged in a row. Instead the hydrogen atoms are mainly on one side of the water molecule. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Representing Molecules The arrangement of atoms within a molecule is called its molecular structure. The molecular structure of ethanol (C2H6O) is more complicated. Each carbon is bonded to four atoms, each hydrogen is bonded to one atom, and the one oxygen is bonded to two atoms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Acetylsalicylic acid, also known as aspirin, has a molecular formula of C9H8O4. What elements make up acetylsalicylic acid? How many atoms of each element are found in one molecule of acetylsalicylic acid? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Acetylsalicylic acid, also known as aspirin, has a molecular formula of C9H8O4. What elements make up acetylsalicylic acid? How many atoms of each element are found in one molecule of acetylsalicylic acid? One molecule of acetylsalicylic acid is made of 9 carbon atoms, 8 hydrogen atoms, and 4 oxygen atoms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Comparing Molecular and Ionic Compounds
What representative units define molecular compounds and ionic compounds? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The representative unit of a molecular compound is a molecule. For an ionic compound, the smallest representative unit is a formula unit. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Recall that a formula unit is the lowest whole-number ratio of ions in an ionic compound. It is important not to confuse formula units with molecules. A molecule is made up of two or more atoms that act as a unit. No such discrete units exist in an ionic compound, which consists of a continuous array of ions. There is no such thing as a molecule of sodium chloride or magnesium chloride. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds. Many molecular compounds are gases or liquids at room temperature. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds. Many molecular compounds are gases or liquids at room temperature. In contrast to ionic compounds, which are formed from a metal combined with a nonmetal, most molecular compounds are composed of atoms of two or more nonmetals. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Water, which is a molecular compound, and sodium chloride, which is an ionic compound, are compared here. Array of sodium ions and chloride ions Collection of water molecules Formula unit of sodium chloride Molecule of water Chemical formula H2O NaCl Chemical formula Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Large differences in electronegativity between two bonded atoms favor the transfer of electrons from the less electronegative (more electropositive) atom to the more electronegative atom resulting in a bond between the two atoms that is IONIC. Smaller differences result in a more equitable “sharing” of electrons between the bonded atoms, resulting in a COVALENT bond between the two atoms. The kinds of bonds formed between elements (covalent vs ionic) can be determined by comparing electronegativity of the two elements. Draw the boundaries on your periodic table

8.2 The Nature of Covalent Bonding
Generally: Nonmetallic element + nonmetallic element  Molecular compound. Molecular compounds are typically gases, liquids, or low melting point solids and are characteristically poor conductors. Examples are H2O, CH4, NH3. Octet Rule: In covalent bonds electrons are usually shared so that all atoms obtain the electron configuration of a noble gas Structural Formula: Shows the covalent bonds as dashes between the bonded atoms Covalent Bonds can be single, double or triple. H-H O=O N≡N Coordinate Covalent Bonds occur when one atom provides both of the shared electrons in a bond. :C≡O:

The Octet Rule in Covalent Bonding
What is the result of electron sharing in covalent bonds? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. For example, a single hydrogen atom has one electron. But a pair of hydrogen atoms shares electrons to form a covalent bond in a diatomic hydrogen molecule. Each hydrogen atom thus attains the electron configuration of helium, a noble gas with two electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Combinations of atoms of the nonmetals and metalloids in Groups 4A, 5A, 6A, and 7A of the periodic table are likely to form covalent bonds. The combined atoms usually acquire a total of eight electrons, or an octet, by sharing electrons, so that the octet rule applies. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds The hydrogen atoms in a hydrogen molecule are held together mainly by the attraction of the shared electrons to the positive nuclei. Two atoms held together by sharing one pair of electrons are joined by a single covalent bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds Hydrogen gas consists of diatomic molecules whose atoms share only one pair of electrons, forming a single covalent bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots. The pair of shared electrons forming the covalent bond is also often represented as a dash, as in H—H for hydrogen. A structural formula represents the covalent bonds as dashes and shows the arrangement of covalently bonded atoms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds The halogens also form single covalent bonds in their diatomic molecules. Fluorine is one example. By sharing electrons and forming a single covalent bond, two fluorine atoms each achieve the electron configuration of neon. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds In the F2 molecule, each fluorine atom contributes one electron to complete the octet. Notice that the two fluorine atoms share only one pair of valence electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair or a nonbinding pair. In F2, each fluorine atom has three unshared pairs of electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds As you can see in the electron dot structures below, the oxygen atom in water has two unshared pairs of valence electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds Methane contains four single covalent bonds. The carbon atom has four valence electrons and needs four more valence electrons to attain a noble-gas configuration. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds Each of the four hydrogen atoms contributes one electron to share with the carbon atom, forming four identical carbon–hydrogen bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds When carbon forms bonds with other atoms, it usually forms four bonds, as in methane. You would not predict this pattern based on carbon’s electron configuration, shown below. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Single Covalent Bonds The formation of four bonds by carbon can be explained by the fact that one of carbon’s 2s electrons is promoted to the vacant 2p orbital to form the following electron configuration. This electron promotion requires only a small amount of energy, and the stability of the resulting methane more than compensates for the small energy cost. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Drawing an Electron Dot Structure
Sample Problem 8.1 Drawing an Electron Dot Structure Hydrochloric acid (HCl (aq)) is prepared by dissolving gaseous hydrogen chloride (HCl (g)) in water. Hydrogen chloride is a diatomic molecule with a single covalent bond. Draw the electron dot structure for HCl. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Solve Apply concepts to the problem.
Sample Problem 8.1 Solve Apply concepts to the problem. Draw the electron dot structures for the hydrogen and chlorine atoms. 2 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Solve Apply concepts to the problem.
Sample Problem 8.1 Solve Apply concepts to the problem. Draw the electron dot structure for the hydrogen chloride molecule. 2 Through electron sharing, the hydrogen and chlorine atoms attain the electron configurations of the noble gases helium and argon, respectively. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Double and Triple Covalent Bonds Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two or three pairs of electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Double and Triple Covalent Bonds A double covalent bond is a bond that involves two shared pairs of electrons. Similarly, a bond formed by sharing three pairs of electrons is a triple covalent bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Double and Triple Covalent Bonds The carbon dioxide (CO2) molecule contains two oxygens, each of which shares two electrons with carbon to form a total of two carbon–oxygen double bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Double and Triple Covalent Bonds Nitrogen (N2), a major component of Earth’s atmosphere, contains triple bonds. A single nitrogen atom has five valence electrons; each nitrogen atom in the molecule must share three electrons to have the electron configuration of neon. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Double and Triple Covalent Bonds You might think that an oxygen atom, with six valence electrons, would form a double bond by sharing two of its electrons with another oxygen atom. In such an arrangement, all the electrons within the molecule would be paired. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Double and Triple Covalent Bonds Experimental evidence, however, indicates that two of the electrons in O2 are still unpaired. Thus, the bonding in the oxygen molecule (O2) does not obey the octet rule. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Nitrogen and oxygen are both diatomic molecules; the table below lists some other diatomic molecules. Diatomic Elements Name Chemical formula Electron dot structure Properties and uses Fluorine F2 Greenish-yellow reactive toxic gas. Compounds of fluorine, a halogen, are added to drinking water and toothpaste to promote healthy teeth. Bromine Br2 Dense red-brown liquid with pungent odor. Compounds of bromine, a halogen, are used in the preparation of photographic emulsions. Hydrogen H2 Colorless, odorless, tasteless gas. Hydrogen is the lightest known element. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The “octet” in the octet rule refers to eight of what?

The “octet” in the octet rule refers to eight of what?
Each of the atoms joined by a covalent bond usually acquires eight electrons in its valence shell. Most noble gases have eight valence electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

8.3 Bonding Theories Molecular Orbital Theory: when atoms bond their atomic orbitals overlap producing molecular orbitals that belong to the entire molecule. Molecular orbitals that are occupied by two electrons of a covalent bond are called bonding orbitals. Sigma bonds: the pair of bonded electrons are found along the axis of the bonded nuclei. Pi bonds: the pair of bonded electrons are found above and/or below the axis of the bonded nuclei.

VSEPR – Valence Shell Electron Pair Repulsion theory
Says that bonded electrons in small molecules will arrange themselves as far apart as possible. Used to predict the shape of small molecules

Molecular Geometry - Linear
CO2 Carbon Dioxide Linear molecules All molecules that contain only 2 atoms and some with 3 Bond angles 180º Be & C with sp hybridization O2, CO2, N2, HCl linear 2 bonding regions 0 lone pairs BeCl2

Molecular Geometry – Trigonal Planar
3 bonding regions 0 lone pairs trigonal planar BF3 Trigonal Planar Molecules 1 central atom surrounded by 3 equally spaced atoms in one plane. The central atom has no unbonded pairs of electrons 120º Bond angles Carbon or boron with sp2 hybridization as central atom BCl3, BF3, H2CO

Molecular Geometry – Tetrahedral
Tetrahedral Molecules 1 central atom surrounded by 4 equally spaced atoms in 3-D. The central atom has no unbonded pairs of electrons 109.5º Bond angles Carbon or silicon with sp3 hybridization as the central atom CH4, NH4+, SiCl4 4 bonding regions 0 lone pairs tetrahedral NH4+

Molecular Geometry – Pyramidal
3 bonding regions 1 lone pair trigonal pyramidal NH3 Pyramidal 1 central atom surrounded by 3 bonded atoms and one unbonded pair of electrons. The unbonded pair of electrons repel the bonded pairs into a 3-D pyramid 107º Bond angles Nitrogen or Phosphorus as the central atom NH3, PCl3, NH3 ammonia

Molecular Geometry – Bent
2 bonding regions 2 lone pairs bent NH2- Bent 1 central atom surrounded by 2 bonded atoms and two unbonded pairs of electrons. The unbonded pairs of electrons repel the bonded pairs into a bent molecule 105º Bond angles Oxygen or Sulfur as the central atom H2O, SO2, NH2- H2O water

Hybrid Orbitals The changing or combining of orbitals within an atom which creates new bond angles. sp – linear - Be & C sp2 – trigonal planar – B & C sp3 – tetrahedral – C & Si

Bond Length Br-Br N-N 145 pm N=N 125 pm NºN 110 pm Cl-Cl
The greater the number of bonds the shorter the bond length N-N pm N=N pm NºN pm The larger the atoms the longer the bond length F-F 142 pm Cl-Cl 199 pm Br-Br 228 pm

Molecular Orbitals Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole. A molecular orbital that can be occupied by two electrons of a covalent bond is called a bonding orbital. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals Sigma Bonds When two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei, a sigma bond is formed. Its symbol is the Greek letter sigma (σ). s atomic orbital Bond axis Sigma-bonding molecular orbital  represents the nucleus Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals Sigma Bonds
In general, covalent bonding results from an imbalance between the attractions and repulsions of the nuclei and electrons involved. The nuclei and electrons attract each other. Nuclei repel other nuclei. Electrons repel other electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

In a bonding molecular orbital of hydrogen, however, the attractions between the hydrogen nuclei and the electrons are stronger than the repulsions. The balance of all the interactions between the hydrogen atoms is thus tipped in favor of holding the atoms together. The result is a stable diatomic molecule of H2. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Atomic p orbitals can also overlap to form molecular orbitals. A fluorine atom, for example, has a half-filled 2p orbital. When two fluorine atoms combine, the p orbitals overlap to produce a bonding molecular orbital. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

There is a high probability of finding a pair of electrons between the positively charged nuclei of the two fluorines. The overlap of the 2p orbitals produces a bonding molecular orbital that is symmetrical when viewed around the F—F bond axis connecting the nuclei. Therefore, the F—F bond is a sigma bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals Pi Bonds
In the sigma bond of the fluorine molecule, the p atomic orbitals overlap end to end. In some molecules, however, orbitals can overlap side to side. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

As shown here, the side-by-side overlap of atomic p orbitals produces what are called pi molecular orbitals. p atomic orbital Pi-bonding molecular orbital  represents the nucleus Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

In a pi bond (symbolized by the Greek letter ), the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms. Because atomic orbitals in pi bonding overlap less than in sigma bonding, pi bonds tend to be weaker than sigma bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

How can a drawing show where an electron is most likely to be found?
CHEMISTRY & YOU How can a drawing show where an electron is most likely to be found? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

How can a drawing show where an electron is most likely to be found?
CHEMISTRY & YOU How can a drawing show where an electron is most likely to be found? Drawings can show molecular orbitals, which are the areas where bonding electrons are most likely to be found. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

In molecules with sigma bonds, where are the shared electrons most likely to be found?

In molecules with sigma bonds, where are the shared electrons most likely to be found?
In sigma bonds, bonding electrons are most likely found between the positively charged nuclei of the atoms bonded together. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Coordinate Covalent Bonds
How are coordinate covalent bonds different from other covalent bonds? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Carbon monoxide (CO) is an example of a type of covalent bonding different from that seen in water, ammonia, methane, and carbon dioxide. It is possible for both carbon (which needs to gain four electrons) and oxygen (which needs to gain two electrons) to achieve noble-gas electron configurations by a type of bonding called coordinate covalent bonding. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Look at the double covalent bond between carbon and oxygen. With the double bond in place, the oxygen had a stable electron configuration, but the carbon does not. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

As shown below, the dilemma is solved if the oxygen also donates one of its unshared pairs of electrons for bonding. A covalent bond in which one atom contributes both bonding electrons is a coordinate covalent bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The ammonium ion (NH4+) consists of atoms joined by covalent bonds, including a coordinate covalent bond. A polyatomic ion, such as NH4+, is a tightly bound group of atoms that has a positive or negative charge and behaves as a unit. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The ammonium ion forms when a positively charged hydrogen ion (H+) attaches to the unshared electron pair of an ammonia molecule (NH3). Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Most polyatomic cations and anions contain covalent and coordinate covalent bonds. Therefore, compounds containing polyatomic ions include both ionic and covalent bonding. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Drawing the Electron Dot Structure of a Polyatomic Ion
Sample Problem 8.2 Drawing the Electron Dot Structure of a Polyatomic Ion The H3O+ ion forms when a hydrogen ion is attracted to an unshared electron pair in a water molecule. Draw the electron dot structure of the hydronium ion. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Analyze Identify the relevant concepts.
Sample Problem 8.2 Analyze Identify the relevant concepts. Each atom must share electrons to satisfy the octet rule. 1 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Solve Apply the concepts to the problem.
Sample Problem 8.2 2 Solve Apply the concepts to the problem. Draw the electron dot structure of the water molecule and the hydrogen ion. Then draw the electron dot structure of the hydronium ion. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sample Problem 8.2 2 Solve Apply the concepts to the problem. The oxygen must share a pair of electrons with the added hydrogen ion to form a coordinate covalent bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sample Problem 8.2 Solve Apply the concepts to the problem. 2 Check that all the atoms have the electrons they need and that the charge is correct. The oxygen in the hydronium ion has eight valence electrons, and each hydrogen shares two valence electrons, satisfying the octet rule. The water molecule is neutral, and the hydrogen ion has a positive charge, giving the hydronium ion a charge of 1+. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Do all atoms joined in covalent bonds donate electrons to the bond?

Do all atoms joined in covalent bonds donate electrons to the bond?
No. In coordinate covalent bonds, the shared electron pair comes from one of the bonding atoms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Exceptions to the Octet Rule
What are some exceptions to the octet rule? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. There are also molecules in which an atom has less, or more, than a complete octet of valence electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Two plausible electron dot structures can be drawn for the NO2 molecule, which has a total of seventeen valence electrons. It is impossible to draw an electron dot structure for NO2 that satisfies the octet rule for all atoms, yet NO2 does exist as a stable molecule. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Some molecules with an even number of valence electrons, such as some compounds of boron, also fail to follow the octet rule. A few atoms, especially phosphorus and sulfur, expand the octet to ten or twelve electrons. Sulfur hexafluoride (SF6) is an example. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Are molecules that do not obey the octet rule necessarily unstable?

Are molecules that do not obey the octet rule necessarily unstable?
No. There are molecules like NO2 that do not obey the octet rule, but that are stable, naturally occurring molecules. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

A large quantity of heat is released when hydrogen atoms combine to form hydrogen molecules. This release of heat suggests that the product is more stable than the reactants. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The covalent bond in the hydrogen molecule (H2) is so strong that it would take 435 kJ of energy to break apart all of the bonds in 1 mole (about 2 grams) of H2. The energy required to break the bond between two covalently bonded atoms is known as the bond dissociation energy. The units for this energy are often given in kJ/mol, which is the energy needed to break one mole of bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

A large bond dissociation energy corresponds to a strong covalent bond. A typical carbon–carbon single bond has a bond dissociation energy of 347 kJ/mol. Strong carbon–carbon bonds help explain the stability of carbon compounds. They are unreactive partly because the dissociation energy is high. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

True or False: A strong covalent bond has a low bond dissociation energy.

True or False: A strong covalent bond has a low bond dissociation energy.
False. A large bond dissociation energy corresponds to a strong covalent bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Resonance Resonance How are resonance structures used?

Resonance The ozone molecule has two possible electron dot structures.
Notice that the structure on the left can be converted to the one on the right by shifting electron pairs without changing the positions of the oxygen atoms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Resonance Because earlier chemists imagined that the electron pairs rapidly flip back and forth, or resonate, between the different electron dot structures, they used double-headed arrows to indicate that two or more structures are in resonance. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Resonance Double covalent bonds are usually shorter than single bonds, so it was believed that the bond lengths in ozone were unequal. Experimental measurements show, however, that the two bonds in ozone are the same length. The actual bonding is a hybrid, or mixture, of the extremes represented by the resonance forms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Resonance The two electron dot structures for ozone are examples of what are still referred to as resonance structures. Resonance structures are structures that occur when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Resonance Chemists use resonance structures to envision the bonding in molecules that cannot be adequately described by a single structural formula. Although no back-and-forth changes occur, double-headed arrows are used to connect resonance structures. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Do resonance structures accurately represent actual bonding?

Do resonance structures accurately represent actual bonding?
No. Resonance structures are a way to envision the bonding in certain molecules. The actual bonding is a hybrid, or mixture, of the extremes represented by the resonance forms. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

8.4 Polar Bonds and Molecules
Polar molecules have oppositely charged ends (dipoles) Two factors determine polarity Type of bonds – Polar or Non-polar Geometry of the molecule (symmetry) Symmetrical molecules are non-polar Nonsymmetrical molecules are polar if they have polar bonds and non-polar if they have non-polar bonds. Large molecules have combinations of the basic shapes that are bent and twisted by polar regions DNA, Proteins, Enzymes Water is polar due to its bent shape so it has unique properties Universal solvent – dissolves polar substances and small molecules Liquid at room temperature Solid less dense than liquid

Sample Problem 8.3 Identifying Bond Type Which type of bond (nonpolar covalent, moderately polar covalent, very polar covalent, or ionic) will form between each of the following pairs of atoms? a. N and H b. F and F c. Ca and Cl d. Al and Cl Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Analyze Identify the relevant concepts.
Sample Problem 8.3 Analyze Identify the relevant concepts. 1 In each case, the pairs of atoms involved in the bonding pair are given. The types of bonds depend on the electronegativity differences between the bonding elements. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Solve Apply concepts to this problem.
Sample Problem 8.3 Solve Apply concepts to this problem. 2 Identify the electronegativities of each atom using Table 6.2. a. N(3.0), H(2.1) b. F(4.0), F(4.0) c. Ca(1.0), Cl(3.0) d. Al(1.5), Cl(3.0) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sample Problem 8.3 Solve Apply concepts to this problem. 2 Calculate the electronegativity difference between the two atoms. The electronegativity difference between two atoms is expressed as the absolute value. So, you will never express the difference as a negative number. a. N(3.0), H(2.1); 0.9 b. F(4.0), F(4.0); 0.0 c. Ca(1.0), Cl(3.0); 2.0 d. Al(1.5), Cl(3.0); 1.5 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sample Problem 8.3 Solve Apply concepts to this problem. 2 Based on the electronegativity difference, determine the bond type using Table 8.4. a. N(3.0), H(2.1); 0.9; moderately polar covalent b. F(4.0), F(4.0); 0.0; nonpolar covalent c. Ca(1.0), Cl(3.0); 2.0; ionic d. Al(1.5), Cl(3.0); 1.5; very polar covalent Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Describing Polar Covalent Molecules
The presence of a polar bond in a molecule often makes the entire molecule polar. In a polar molecule, one end of the molecule is slightly negative, and the other end is slightly positive. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

In the hydrogen chloride molecule, for example, the partial charges on the hydrogen and chlorine atoms are electrically charged regions, or poles. A molecule that has two poles is called a dipolar molecule, or dipole. The hydrogen chloride molecule is a dipole. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds. A carbon dioxide molecule has two polar bonds and is linear. O C O Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The water molecule also has two polar bonds. However, the water molecule is bent rather than linear. Therefore, the bond polarities do not cancel and a water molecule is polar. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

What is the difference between an ionic bond and a very polar covalent bond?

What is the difference between an ionic bond and a very polar covalent bond?
Two atoms will form an ionic bond rather than a very polar covalent bond if the two atoms have a slightly higher difference in electronegativity—a difference of more than 2.0. There is no sharp boundary between ionic and covalent bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecules can be attracted to each other by a variety of different forces. Intermolecular attractions are weaker than either ionic or covalent bonds. Among other things, these attractions are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces The two weakest attractions between molecules are collectively called van der Waals forces, named after the Dutch chemist Johannes van der Waals. Van der Waals forces consist of dipole interactions and dispersion forces. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Dipole interactions occur when polar molecules are attracted to one another. The electrical attraction occurs between the oppositely charged regions of polar molecules. Dipole interactions are similar to, but much weaker than, ionic bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces The slightly negative region of a polar molecule is weakly attracted to the slightly positive region of another polar molecule. Dipole interactions are similar to, but much weaker than, ionic bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Dispersion forces, the weakest of all molecular interactions, are caused by the motion of electrons. They occur even between nonpolar molecules. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Dispersion forces, the weakest of all molecular interactions, are caused by the motion of electrons. When the moving electrons happen to be momentarily more on the side of a molecule closest to a neighboring molecule, their electric force influences the neighboring molecule’s electrons to be momentarily more on the opposite side. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Dispersion forces, the weakest of all molecular interactions, are caused by the motion of electrons. The strength of dispersion forces generally increases as the number of electrons in a molecule increases. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Fluorine and chlorine have relatively few electrons and are gases at ordinary room temperature and pressure because of their especially weak dispersion forces. Bromine molecules therefore attract each other sufficiently to make bromine a liquid under ordinary room temperature and pressure. Iodine, with a still larger number of electrons, is a solid at ordinary room temperature and pressure. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hydrogen Bonds The dipole interactions in water produce an attraction between water molecules. Each O—H bond in the water molecule is highly polar, and the oxygen acquires a slightly negative charge because of its greater electronegativity. The hydrogens in water molecules acquire a slightly positive charge. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hydrogen Bonds This relatively strong attraction, which is also found in hydrogen-containing molecules other than water, is called a hydrogen bond. Hydrogen bond Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hydrogen Bonds Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom. The other atom may be in the same molecule or in a nearby molecule. Hydrogen bonding always involves hydrogen. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hydrogen Bonds A hydrogen bond has about 5 percent of the strength of the average covalent bond. Hydrogen bonds are the strongest of the intermolecular forces. They are extremely important in determining the properties of water and biological molecules. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Intermolecular forces between molecules are weaker than ionic or covalent bonds. Van der Waals Forces: generally weak attractions between nonpolar molecules caused by induced dipoles on the attracted molecules. Also called London Dispersion Force (LDF). Dipole-Dipole forces: occur between the positive and negative ends of adjacent polar molecules. Hydrogen Bonds: attractive forces in which a hydrogen atom is covalently bonded to a highly electronegative element (F,O,Cl, or N) and is attracted to an unbonded pair of electrons on another atom or molecule. This results in a strong intermolecular force (5% of covalent bond strength). Network solids: molecular substances with very strong forces between atoms/molecules of the same type (diamond/quartz).

Why are hydrogen bonds important?

Why are hydrogen bonds important?
Hydrogen bonds are the strongest of the intermolecular forces and are extremely important in determining the properties of water and biological molecules such as proteins. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

At room temperature, some compounds are gases, some are liquids, and some are solids. The physical properties of a compound depend on the type of bonding it displays—in particular, on whether it is ionic or covalent. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The diversity of physical properties among covalent compounds is mainly because of widely varying intermolecular attractions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The melting and boiling points of most compounds composed of molecules are low compared with those of ionic compounds. In most solids formed by molecules, only the weak attractions between molecules need to be broken. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

A few solids that consist of molecules do not melt until the temperature reaches 1000°C or higher. Most of these very stable substances are network solids (or network crystals), solids in which all of the atoms are covalently bonded to each other. Melting a network solid would require breaking covalent bonds throughout the solid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Diamond is an example of a network solid. Each carbon atom in a diamond is covalently bonded to four other carbons, interconnecting carbon atoms throughout the diamond. Diamond does not melt; rather, it vaporizes to a gas at 3500°C and above. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

This table summarizes some of the characteristic differences between ionic and covalent (molecular) substances. Characteristics of Ionic and Molecular Compounds Characteristic Ionic Compound Molecular Compound Representative unit Formula unit Molecule Bond formation Transfer of one or more electrons between atoms Sharing of electron parts between atoms Type of elements Metallic and nonmetallic Nonmetallic Physical state Solid Solid, liquid, or gas Melting point High (usually above 300°C) High (usually below 300°C) Solubility in water Usually high High to low Electrical conductivity of aqueous solution Good conductor Poor to nonconducting Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Why do network solids take so much more heat to melt than most covalent compounds?

Why do network solids take so much more heat to melt than most covalent compounds?
Melting a network solid requires breaking covalent bonds throughout the solid. Melting most covalent compounds only requires breaking the weak attractions between molecules. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

NAMING COMPOUNDS The chemical formula represents the composition of each molecule. In writing the chemical formula, in almost all cases the element farthest to the left (the most metallic) in the periodic table is written first. So for example the chemical formula of a compound that contains one sulfur atom and six fluorine atoms is SF6. If the two elements are in the same period, the symbol of the element of that is lower in the group (less electronegative) is written first e.g. IF3.

9.2 NAMING BINARY IONIC COMPOUNDS
contain only 2 elements a metal cation and a nonmetal anion Compounds formed by elements on opposite sides of the periodic table which either give up (left side) or take up electrons (right side). Depending on the atom, there can be an exchange of more than one electron resulting in charges greater than ±1. Always write the name of the metallic cation first followed by the anion and end in “ide”.

Group IA – alkali metals – loose 1 e- to form +1 (Na+)
Group II A– alkaline earth metals –loose 2 e- to form +2 (Ca+2) Group III A– loose three e- to form +3 (Al+3) Group IV A– loose four e- to form +4 (Sn+4) Group V A– accept three e- to form –3 (N-3) Group VI A– accept two e- to form –2 (O-2) Group VIIA – accept one e- to form –1 (Cl-1)

Naming IONIC compounds
Anions – suffix – “ide” So Cl- is chloride Oxygen O2- is OXIDE S2- is SULFIDE Cations For Na+, Ca2+, the name of the ion is the same except refer to the ion. So SODIUM ION or SODIUM CATION NaCl - sodium chloride CaCl2 - calcium chloride

NAMING TERNARY COMPOUNDS
Contain polyatomic ions – more than 2 elements present Give the name of the cation followed by the name of the anion. Positive Molecular Ions End the name with “ium” or “onium” NH4+ is ammonium, H3O+ is hydronium Negative Molecular Ions – use the list of polyatomic ions - NO3 - NITRATE 2- SO4 - SULFATE NO2 - NITRITE 3- PO4 - PHOSPHATE

Transition Elements The transition elements are chemically quite different from the metals in the “A” block, due to differences in electronic configuration For example, Fe can loose two or three electrons to become Fe2+ and Fe3+, respectively.

STOCK SYSTEM for naming ionic compounds
To identify the charge of Fe in a compound the following nomenclature is used. A Roman numeral indicates the charge on the cation. Fe2+ is iron(II) Fe3+ is iron (III) So iron(III) chloride is FeCl3 TRADITIONAL SYSTEM for naming ionic compounds An older scheme differentiated between the lower and higher charge by ending the name of the element with “ous” to indicate the lower charge and “ic” for the higher. ferrous chloride => FeCl2 ferric chloride => FeCl3 However, this convention does not indicate the numerical value of the charge.

Traditional System for Molecular compounds
9.3 Naming Molecular Compounds: the name of the first element in the formula is unchanged. The suffix “-ide” is added to the second element. Traditional System for Molecular compounds Add a prefix to the name of the second element (and the first element if there is more than one atom) to indicate the number of atoms of that element in the compound mono – 1 hexa – 6 di – 2 hepta – 7 tri – 3 octa – 8 tetra – 4 nona – 9 penta – 5 deca - 10 P4O10 – tetraphosphorous decoxide CO – carbon monoxide SF6 – sulfur hexafluoride CO2 – carbon dioxide

Stock System for Molecular Compounds
“Pretend” the compound is ionic and give the name of the more metallic element first followed by a Roman Numeral to indicate the “apparent” charge on that element. Assume the more electronegative element forms its most common ion. Name Formula Nitrogen II oxide NO Nitrogen V oxide N2O5 Carbon IV Chloride CCl4

9.4 NAMING ACIDS Hydrogen forms binary compounds with almost all non-metals except the noble gases. The binary compounds of hydrogen are special cases. They form acids in water so they have their own nomenclature. Binary Acids - contain hydrogen and one other element. Use the system Hydro____ ic acid and fill in the name of the non hydrogen element HF - hydrogen fluoride - Hydrofluoric Acid HCl - hydrogen chloride – Hydrochloric Acid H2S - hydrogen sulfide – Hydrosulfuric Acid H3P - hydrogen phosphide – Hydrophosphoric Acid

Ternary Acids - contain hydrogen and a polyatomic ion
Ternary Acids - contain hydrogen and a polyatomic ion. Give the name of the ion and end in “-ic” or “-ous”. “-ate” ions become “-ic” acids. “-ite” ions become “-ous” acids. “-ate” Ion Predicted Name Acid Name HNO3 hydrogen nitrate Nitric Acid H2SO4 hydrogen sulfate Sulfuric Acid H3PO4 hydrogen phosphate Phosphoric Acid H2CO3 hydrogen carbonate Carbonic Acid HC2H3O hydrogen acetate Acetic Acid “-ite” Ion Predicted Name Acid Name HNO2 hydrogen nitrite Nitrous Acid H2SO3 hydrogen sulfite Sulfurous Acid H3PO3 hydrogen phosphite Phosphorous Acid

Organic molecules (containing C) have a separate nomenclature
The molecular formulas for compounds containing C and H (called hydrocarbons) are written with C first. Example, CH4, C2H6, etc. Use a prefix for the number of carbon atoms 1-meth 2-eth 3-prop 4-but 5-pent 6-hex 7-hept 8-oct 9-non 10-deca Single bonds between C atoms suffix “ane”. Double bonds between C atoms suffix “ene” Triple bonds between C atoms suffix “yne” Law of multiple proportions: the ratio of masses of an element in 2 compounds that contain the same elements is a simple whole number ratio.

Chemistry I Notes Unit 3 Chapters 7-9

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