UNIT 2 Chemistry in Action.

UNIT 2 Chemistry in Action

Specific Learning Outcomes
Relate an element’s position on the periodic table to its combining capacity (valence). Include: alkali metals, alkaline earths, chalcogens, halogens, noble gases. Explain, using the periodic table, how and why elements combine in specific ratios to form compounds. Include: ionic bonds, covalent bonds. Write formulas and names of binary ionic compounds. Include: IUPAC guidelines and rationale for their use.

Write formulas and names for covalent compounds using prefixes
Write formulas and names for covalent compounds using prefixes. Include: mono, di, tri, tetra. Investigate the Law of Conservation of Mass, and recognize that mass is conserved in chemical reactions. Balance chemical equations.

Investigate and classify chemical reactions as synthesis, decomposition, single displacement, double displacement, or combustion. Perform an experiment to classify acids and bases using their characteristic properties. Include: pH, indicators, reactivity with metals. Explain how acids and bases interact to form a salt and water in the process of neutralization.

Review of the Atom Atoms can be broken down into subatomic particles. The three main subatomic particles are protons, neutrons, and electrons. Nucleus If an atom were the size of a football stadium, the nucleus would be the size of a football. • A nucleus is made of protons and neutrons. — Protons and neutrons both have a mass of 1μ. — Protons have a positive charge (+1); neutrons have no charge. — The total mass of the atom is equal to the sum of the masses of the protons and neutrons in the nucleus. • If an atom were the size of a football stadium, the nucleus would be the size of a football. • The atomic number of an atom is equal to the number of protons in the nucleus. • The atomic mass number is equal to the sum of the number of protons and neutrons. The atomic mass is expressed in atomic mass units (μ). The modern atomic theory states that an atom consists of: • Protons located in the nucleus of the atom. Protons have a positive charge and have a mass of one atomic mass unit. An atomic mass unit is defined as l/12th the mass of a carbon atom that has six protons and six neutrons. This means that 1μ p is equal to the mass of a proton. • Neutrons also located in the nucleus of the atom. Neutrons have no electrical charge and have a mass of approximately 1μ. Electrons • Electrons located around the nucleus in less well-defined orbits than first thought. Electrons have a single negative electrical charge but their mass is considered zero since it is so small (approximately 1/2000 the mass of a proton). • Electrons move around the nucleus in specific paths called energy levels. — Energy levels exist whether there is an electron in them or not. — Electrons occupy certain energy levels depending on the atom. For example, hydrogen has a single electron in the first energy level; sodium has two electrons in the first energy level, eight electrons in the second energy level, and one electron in the third energy level. — An atom can have a maximum of two electrons occupying the first energy level, eight in the second energy level, and eight in the third energy level. We will not study the structure of atoms with more than three levels of energy levels. Each energy level must be filled before electrons occupy the next one.

Charge Location Mass Protons +1 Nucleus 1 amu Electrons - Outer shells
Negligible Neutrons No. of electrons = no. of protons Atomic number = no. of protons Mass number = no. of protons and neutrons amu = atomic mass unit

Determining the Number of Subatomic particles
Number of protons The number of protons is equal to the atomic number. Number of electrons In any atom, the number of electrons is equal to the atomic number. Number of neutrons The number of neutrons is equal to the mass number (rounded to a whole number) minus the atomic number.

Atomic number Atomic mass

Protons = 6 Electrons = 6 Neutrons = 12 – = 6

Mass number = atomic number (number of protons) + number of neutrons
Atomic number = mass number – number neutrons

Bohr’s Atomic Model Bohr's model of the atom proposed that electrons occupy orbits or energy levels or shells. Bohr discovered that the location of each shell was a certain distance from the nucleus. He also discovered that only a specific number of electrons populated each shell. Hydrogen has one electron in the first energy level, and helium, the next element, has two electrons. The inner energy level cannot hold more than two electrons, so helium's inner energy level is considered full. Lithium has three electrons. The first two electrons occupy the inner energy level and the third occupies the second energy level. The second orbit is eventually filled when it has eight electrons. Neon follows fluorine and has 10 electrons in its orbits, a) How many electrons will neon have in its first orbit? 2 b) How many electrons will neon have in its second orbit? 8 c) The next element, sodium, has one more electron in its orbit. Where will that electron be located? In the 3rd orbit

When electrons occupy the shells, they begin at the closest shell to the nucleus. The first shell can contain a maximum of two electrons, and the second and third shells allow a maximum of eight electrons.

Nucleus Protons 1 p+ Electrons 1 e-

Nucleus First orbit (max 2e- ) Second orbit (max 8e- ) Third orbit (max 8e- )

Lithium Nucleus The third electron in Li occupies the second energy level (shell)

Fluorine The first 2 electrons are located in the first shell.
The next 7 electrons occupy the second shell.

When an element has more than four electrons in its shell, you arrange the dots in twos.
These electrons are important in helping to determine how an element will react with other materials. These electrons in the outermost energy level are called valence electrons.

The Periodic Table Dmitri Mendeleev, a Russian scientist and professor, placed characteristic properties of elements on pieces of paper and arranged them in many different ways. He discovered a certain pattern or repetition of properties. This repetition of properties or periodicity was an important outcome of Mendeleev's work. Mendeleev used the periodicity of properties to create a periodic table. This table summarized the structure and properties of the elements.

Today's periodic law states that the properties of elements are a periodic function of their atomic numbers. The periodic table is an important tool for chemists. It quickly determines some key facts about an element. The basic information with regard to the structure of an atom is quickly established with the periodic table.

Al 13 +3 26.98 Chemistry : Periodic Table & Elements
The basic information with regard to the structure of an atom is quickly established with the periodic table. 26.98 Chemistry : Periodic Table & Elements

Columns in the Periodic Table
The periodic table arranges the elements in columns. A single column is called a group or family. A family contains elements that have similar but not identical properties.

Alkali Metals (H, Li, Na, K, Rb, Cs, Fr)
Occupies the first column in the periodic table. Has one valence electron in its outer energy level. Most reactive metals In their natural state, alkali metals are always found combined with other substances because of their reactivity. The alkali metal family occupies the first column in the periodic table. I t includes lithium (Li), sodium (Na), potassium (K), etcetera. Each element has one valence electron in its outer energy level. These metals are the most reactive metals in the periodic table because of the single electron in the outer energy level. In their natural state, alkali metals are always found combined with other substances because of their reactivity. The most common element in the family is sodium, which is found all over the Earth in compounds like salt (sodium chloride).

Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)
Located in the second column of the periodic table. Less reactive than the alkaline metal family. Has two valence electrons in the outer energy level. The alkaline earth metals family is located in the second column of the periodic table. Alkaline earth metals are less reactive than the alkaline metal family. Their lesser activity arises from having two valence electrons in the outer energy level. Beryllium is the first member of the family, followed by magnesium, calcium, strontium, barium, and radium.

Chalcogens (O, S, Se, Te, Po)
Located in the 16th column. Slightly less reactive than the halogen family. Have six valence electrons in their outer energy level The chalcogen (oxygen) family is located in the 16th column of the periodic table. The chalcogen family is slightly less reactive than the halogen family. They have six valence electrons in their outer energy level (that is, they are two electrons short of having a completely filled outer energy level). The first member of the chalcogen family is oxygen, followed by sulphur, selenium, tellurium, and polonium.

Halogens (F, Cl, Br, I, At) The halogen family is the 17th family
Halogens such as fluorine and chlorine react with one atom of hydrogen to form HF and HCI respectively. Have seven valence electrons Most reactive non-metals. In their natural state, the highly reactive halogens are found combined with another element. The halogen (fluorine) family is the 17t h family in the periodic table and includes fluorine (F), chlorine (CI), bromine (Br), iodine (I), and astatine (At). Halogens such as fluorine and chlorine react with one atom of hydrogen to form HF and HCI respectively. The halogens have seven valence electrons (that is, they are one electron short of filling their outermost energy level). The halogens are the most reactive non-metals in the periodic table. In their natural state, the highly reactive halogens are found combined with another element.

Noble Gases (He, Ne, Ar, Kr, Xe, Rn)
The noble gas family is the 18th family Called noble gases because they do not generally form compounds with other elements. Inert, unreactive because their outer energy levels are completely filled with electrons. No natural compounds formed from these gases exist. The noble gases (helium) family is the 18t h family in the periodic table. I t includes helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Ra). They are called noble gases because they do not generally form compounds with other elements. They are unreactive because their outer energy levels are completely filled with electrons. No natural compounds formed from these gases exist.

Hydrogen Sometimes hydrogen behaves as a metal and sometimes as a non-metal. Hydrogen has one electron in its outermost energy level, so it is reactive. Almost all the hydrogen on Earth is combined with other materials or with itself. Hydrogen is a special case because it is a family of one. Sometimes hydrogen behaves as a metal and sometimes as a non-metal. Hydrogen has one electron in its outermost energy level, so it is reactive. Almost all the hydrogen on Earth is combined with other materials or with itself.

You should be able to predict with some accuracy the properties of elements that are in the same families.

Look at your periodic table and note the number of valence electrons in each family. State the number of valence electrons for the families in the table shown below.

Number of Family Valence Electrons alkali metals 1
alkaline earth metals 2 chalcogens 6 halogens 7 noble gases 8

Rows in the Periodic Table
Rows in the periodic table are called periods. Elements in periods do not demonstrate similar properties as they do in families. Periods, however, show trends. Periods each represent a new energy level filled up with electrons. Video: How to Use the Periodic Table Video: “The Elements”

Electron Dot Diagram An electron dot diagram or Lewis diagram represent an atom and its valence electrons. The electrons in the valence shell are shown as dots placed around the symbol. Electron dot diagrams are a valuable tool for describing, predicting, and explaining compound formation. -Shows how valence electrons are arranged among atoms, ions, molecules and compounds. -Valence electrons are electrons in the outermost shell of an atom. The group number is equal to the number of valence electrons.

Ca Example: Draw an electron dot diagram for calcium.
Calcium #20, Group II

S Example: Draw an electron dot diagram for sulphur.
Sulfur, #16, Group VI

N Example: Draw an electron dot diagram for nitrogen.
Nitrogen, #7, Group V Video: Valence Electrons & Lewis Diagrams Video: How to draw Lewis Diagrams

Ionic Bonds Ionic bonds result when electrons are transferred from metal atoms to non-metal atoms. The metal atoms lose electrons to become positive ions, while the non-metal atoms gain electrons to become negative ions. The ions are then held together by the action of opposite charges in an ionic bond. Atoms have a tendency to combine and form new materials. Only approximately 100 different atoms combine to form the millions of materials that are found on Earth. In some ways, atoms behave like letters of the alphabet in that 26 letters combine to form the huge number of words in the English language. An atom has gained or lost one or more electrons No atom, aside from the elements in group 8, has a full set of electrons in its outer shell An ion’s trick is to transfer electrons to and from the outer shells to fill them up. This creates positively and negatively charged ions, making chemical reactions possible. Groups (metals) lose electrons Groups (non metals) gain electrons

Atoms are considered neutral because the number of protons in the nucleus and the number of electrons in the shells around the nucleus are equal, resulting in a zero net charge for the atom. Whenever an atom is discussed, you may assume it is electrically neutral. Atoms contain electric charges—positively charged protons and negatively charged electrons. Atoms are considered neutral because the number of protons in the nucleus and the number of electrons in the energy levels around the nucleus are equal, resulting in a zero net charge for the atom.

Cations (Positive Ions)
What happens if an electron is removed from an atom? If an electron is removed from sodium, the number of positive charges and negative charges are no longer balanced. Sodium has 11 protons and 11 electrons in a neutral atom; but there are 11 protons and 10 electrons when an electron is removed, so the atom has a net positive charge.

11 11 11 10 11 – 10 = +1 Sodium atom Sodium ion
How many positive charges in the atom? How many positive charges in the ion? How many negative charges in the atom? How many negative charges in the ion? What is the residual (net) charge in the atom? What is the residual (net) charge in the ion? 11 11 Sodium is an alkali metal. It is located in the first family and second period of the periodic table. Sodium has a single electron in its outer energy level since it belongs to the alkali family. When sodium combines with a non-metal to form a compound, it will lose one electron. If an electron is removed from sodium, the number of positive charges and negative charges is no longer balanced. 11 10 11 – 10 = +1

When an atom from the alkali family reacts with an atom from another element, it will give off its valence electron to the other atom. In giving away its electron, the alkali metal atom has a filled outer shell. Also, in giving away its electron, the atom becomes positively charged with a 1+ charge.

You do not have to write the 1
Na loses this electron Na Na

Atoms from the alkaline earth family have two valence electrons
Atoms from the alkaline earth family have two valence electrons. These atoms will give off the two valence electrons when combining with an atom from another element. After giving away the two valence electrons, the alkaline earth metal atom has a 2+ charge.

Mg loses these electrons
2+ Mg To get charge: 0 – (-2) = +2 Electrons lost

Alkali and alkaline earth metals form positive ions when forming an ionic bond with another element.
Always write the charge with the number first, followed by the sign ie.) 1+

Anions (Negative Ions)
The oxygen family and the halogen family behave differently. These two families readily accept electrons to fill their valence shells so that their electron configuration also resembles a noble gas. The halogens need only one electron to fill their valence shells, so they accept only one electron. Once the extra electron is accepted, a 1- ion is formed.

Oxygen To get charge: 0 + (-2) = -2 8 p
As 2 electrons are placed into the valence shell, the oxygen atom becomes a 2- ion.

Oxygen has a combining capacity of two
Oxygen has a combining capacity of two. The chalcogen family all have the same combining capacity.

Fluorine To get charge: 0 + (-1) = -1 9 p
As the electron is placed into the valence shell, the fluorine atom becomes a 1- ion.

Fluorine has a combining capacity of one
Fluorine has a combining capacity of one. The halogen family all have the same combining capacity.

When a positively charged ion comes near a negatively charged ion, they attract each other and form a bond called an ionic bond. An ionic bond will hold the two ions together to form a compound.

The formation of compounds often takes place vigorously when metals and non-metals are placed together. If a sample of sodium metal is placed in a container of chlorine gas, an explosive reaction takes place and the sodium combines with the chlorine to form sodium chloride. The Bohr model below provides an explanation for this reaction.

Electron from Na has been transferred to Cl
Sodium Chlorine 11 p 12 n 17 p 18 n The sodium ion and chlorine ion are attracted because of their opposite charges and stick together to form salt. The "sticking together" is called an ionic bond. As extremely large numbers of sodium and chlorine atoms undergo this chemical reaction, they form a crystal of salt.

An electrostatic attraction holds the ions together
An electrostatic attraction holds the ions together. The cation and anion are attracted to each other. This is what forms an ionic bond. - + 11 p 12 n 17 p 18 n

Notice that atoms when combine to form ionic compounds, they always gain or lose enough electrons to have a valence shell like its closest noble gas neighbour. Ionic Bonds

Structure of NaCl

Metals and Non-metals Metals
Generally have three or fewer valence electrons Alkali family of elements has the strongest metallic properties The next strongest metallic properties are found in the alkaline Earth family. Metallic properties exist but in decreasing amounts through the boron family. Generally, elements that have three or fewer electrons in the outer energy level are called metals. The strongest metallic properties are found in the first or alkaline family of elements. The next strongest metallic properties are found in the alkaline Earth family. Metallic properties exist but in decreasing amounts through the boron family. Metals are conductors of heat and electricity, and are shiny.

Metals and Non-metals Non-metals
Have five or more electrons in the valence shell. Begin at the chalcogen family and end at the noble gas family. Elements that have five or more electrons in the outer energy level are classified as non-metals. Non-metals begin at the chalcogen (oxygen) family and end at the noble gas family. Nonmetals are generally gases or brittle solids at room temperature.

There are exceptions to this general classification and some families have members that behave as both metals and non-metals (e.g., silicon). These elements are called metalloids. Many periodic tables have stair steps across families at the right side. These steps show the dividing line between metals and non-metals. The elements on the dividing line are metalloids.

State Appearance Metals solids at room temp., except for Hg (liquid) shiny lustre Non-metals some gases at room temp. some solids one liquid (bromine) not very shiny Metalloids solids at room temperature can be shiny or dull

Malleability and ductility
Conductivity Malleability and ductility Metals good conductors of heat and electricity malleable ductile Non-metals poor conductors of heat and electricity brittle not ductile Metalloids may conduct electricity poor conductors of heat

Review of Valence Electrons
Family Valence Electrons Electrons Lost/Gained Charge of Ions Alkaline metals 1 Lose 1 1+ Alkaline earth metals 2 Lose 2 2+ Boron 3 Lose 3 3+ Carbon 4 Lose or gain 4 4+/- Nitrogen 5 Gain 3 3- Chalcogen 6 Gain 2 2- Halogen 7 Gain 1 1- Noble gas 8

When ionic compounds are formed, elements with a positive valence number will combine with elements having a negative valence number. In general terms, metals (families 1 and 2) combine with non-metals (families 16 and 17).

Covalent Bonds Many compounds do not form ionic bonds. These compounds contain two or more non-metallic atoms. For example, C02 is made of two different non-metals, carbon and oxygen. These compounds are formed through the sharing of valence electrons. A covalent bond is formed when two or more non-metallic atoms share valence electrons. When two ions form ionic bonds, they transfer one or more electrons from a metal atom to a non-metal atom. As a result of the electron transfer, one ion has a positive charge (loses electrons) and one has a negative charge (gains electrons). A n attraction exists between these ions, forming an ionic bond, holding them together as an ionic compound. Many compounds, however, do not form ionic bonds. These compounds contain two or more non-metallic atoms. For example, C 0 2 is made of two different non-metals, carbon and oxygen. These compounds are formed through the sharing of valence electrons. A covalent bond is formed when two or more non-metallic atoms share valence electrons. A molecule is: -two or more atoms held together by a covalent bond -two or more non-metals combine making electrically neutral, stable substances -rather than transferring/robbing atoms of their electrons, it allows non-metals to cooperate in filling their outer shells

Two hydrogen atoms form a covalent bond by sharing electrons to produce a hydrogen molecule.
A molecule is the smallest unit of a covalent compound. A molecule has different characteristic properties from the atoms that form it.

Each line is a bond which shares 2e-
The Bohr model for hydrogen shown below illustrates a covalent bond. H Each line is a bond which shares 2e- When two atoms of hydrogen come close to each other, the protons attract each other's electrons. The force is not strong enough to cause an electron transfer (ionic bond), but i t is strong enough to force the electrons to travel i n both of the atoms' orbits, spending most of the time in the position shown in the diagram, between the two nuclei. As a result, the two electrons are shared by both atoms. The hydrogen atom at the left "looks" at its orbit and "sees" two electrons; so does the one at the right. By sharing their electrons, both atoms are satisfied they have filled outer orbits (the outer orbits are the same as for helium) and are stable. Covalent bond

See white board for example
The electrons are shared in the outer shells of both atoms. This covalent bond forms a molecule of hydrogen (H2). The two hydrogen atoms form a diatomic molecule. Covalent Bonds See white board for example

A basic rule in chemistry is that an atom with eight electrons in its outer shell is particularly stable. This need for eight electrons in a covalent bond is called the octet rule. Isn’t it Ionic Song Covalent vs Ionic How to Determine Ionic & Covalent Bonds Chemical Bonds Song

Diatomic Molecules H H2 N N2 O O2 F F2 Cl Cl2 Br Br2 I I2
Name of Element Symbol for One Atom of the Element Formula for One Molecule Hydrogen H H2 Nitrogen N N2 Oxygen O O2 Fluorine F F2 Chlorine Cl Cl2 Bromine Br Br2 Iodine I I2 In each of the cases above, the outer energy levels of the atoms are "filled" with electrons. What does it mean for a energy level to be filled? A energy level is filled when it contains all the electrons it can hold in that particular energy level. When an energy level is filled, it has the same number of electrons as an inert gas and becomes itself inert or unreactive.

The elements forming diatomic gases are unstable as single atoms and combine almost instantaneously to form stable molecules. Diatomic molecules are still classified as elements even though they are molecules. Remember that diatomic molecules are made of only one kind of atom.

When two identical atoms link together, the arrangement is called a diatomic molecule. H-O-F-Br-I-N-Cl the Clown will help you remember diatomic molecules. Take a close look at his ear (and the number 2), and remember that all these elements exist in pairs.

Chemical Formulas Chemistry has its own language. Chemists communicate in this language to describe the millions of known compounds. This communication depends on a standard system of naming and writing the formulas for compounds. Chemists formed a group to standardize the system of naming and called themselves the International Union of Physical and Applied Chemists, or IUPAC.

A chemical formula is a shorthand method to represent compounds that uses the elements' symbols and subscripts. The chemical formula gives the following information: The different elements in the compound. The number of atoms of each element in the compound.

* No subscript indicates only 1 atom is present*
element symbols Subscript tells you amount of each element. Water contains: 2 Hydrogens 1 Oxygen H2O subscript * No subscript indicates only 1 atom is present*

Na2SO4 Contains: 2 sodium atoms 1 sulphur atom 4 oxygen atoms
Types of Chemical Formulas

Naming Ionic Compounds
When naming an ionic compound from its formula follow the rules below: The cation (positive ion) is named first, followed by the anion (negative ion). Write the full name of the metallic element (positive ion). Write the name of the non-metallic element (negative ion) and change the ending to "-ide". The periodic table can be used to predict how elements combine to form compounds. You know that an element with a positive ion combines with an element with a negative ion. The compounds you have examined up to this point are formed by two elements and are called binary compounds. Writing chemical formulas for these compounds is straightforward. I t is also easy to name binary compounds, but you need to know some rules. Until the 18th century, no system for naming compounds existed. Some compounds received common names based on their appearance or use (e.g., oil of vitriol and butter of arsenic, water, baking soda). When the same substance was given two names by different countries because of their different languages, it caused confusion. As scientists from many countries compared notes and used each other's work as building blocks to the next discovery, standardization became necessary. As a general rule, the following two rules apply to naming ionic compounds: • Name the positive ion first by writing the full name of the metallic element. • Name the non-metal ion next by dropping the last syllable(s) of the name of the name of the element and adding the suffix "ide.“ For example, reaction between sodium and chlorine forms sodium chloride, and the formula SrS is named strontium sulfide. Compounds: -combine elements into new substances -are produced and broken apart by chemical reactions -most substances on Earth exist in the form of compounds -total number of known compounds: 61 million

Example: Write the name of NaCl. Step 1: Name the first element. Step 2: Name the second element and change the ending to "-ide". The name of the compound is ______________________. Na = sodium Cl = chlorine  chloride sodium chloride

Example: Write the name of Mg3P2. Step 1: Name the first element. Step 2: Name the root of the second element and add "-ide". The name of the compound is _____________________. Mg = magnesium P = phosphorus  phosphide magnesium phosphide

Writing Ionic Formulas
The following must occur, when writing the formula for ionic compounds. The formula must have the cation first, followed by the anion. The sum of the charges of the ions must be zero. That is, the number of positive charges must equal the number of negative charges. You may not change the charge of the ions to make the ion charges equal zero. When writing correct formulas for these compounds, there are several rules to follow. Beryllium and chlorine will be used as examples. • Write the symbol of the metallic element first. I n the example, Be is written first . The non-metal, chlorine, is written second. • Place the ionic charge number (valence number) of one element at the base (as a subscript) of the other element as shown below. Use only the valence number, not the charge (+ o r - ) . Note that the valence number of beryllium is placed as a subscript to CI and the valence number of chlorine has been placed as a subscript to beryllium. Notice also that the beryllium subscript of 1 has been left out. • If a subscript has a value of one, leave it out. Notice that Be has no subscript i n the example. • Reduce the subscripts when necessary by the greatest common factor (e.g., when magnesium combines with sulfur, the formula might appear as Mg2S2 , but it should be reduced to MgS by dividing both subscripts by two).

The “Cross-Over” Method
Write the ions and their charges side by side. Make the number of the charge of one ion the subscript of the other ion (omitting the + or – sign). Remember we do not write the number one as a subscript. Reduce all subscripts to their simplest form, if necessary.

Al3+ O2- Al2O3 Al3+ and O2- Example:
Write the formula for aluminum oxide. Step 1: Write the ions and their charges. Step 2: Make the number of the charge of one ion the subscript of the other ion. Al3+ and O2- Al3+ O2- Al2O3

Ba2+ F1- BaF2 Example: Write the formula for barium fluoride.
The charge on the fluoride ion is 1-. Since IUPAC rules do not write the number one as a subscript, we leave the barium without a subscript.

Write the formula for magnesium chloride.
Example: Write the formula for magnesium chloride. Write the formula for calcium oxide. See white board See white board Naming Ionic Compounds

Polyatomic Ions Some ions are composed of several atoms joined covalently. These are called polyatomic ions (poly = many). The charge for polyatomic ions is for the whole group of atoms not just for the atom written last. DO NOT change the subscripts of polyatomic ions; if you change the subscripts you change the identity of these ions.

When indicating the presence of more than one polyatomic ion in a compound, we use parenthesis around the polyatomic ion, followed by its subscript. For example, the compound Al(CH3COO)3 has an aluminum ion and 3 acetate ions. Placing the acetate ion in parenthesis and following it with the subscript 3 indicates there are 3 acetate ions.

Example: Write the name for KNO3. Step 1: Identify the cation. Step 2: Identify the anion. Step 3: Write the name of the cation first, followed by the anion. K+  potassium ion NO3-  nitrate ion potassium nitrate

Example: Write the name of Na3PO4. Step 1: Identify the cation. Step 2: Identify the anion. Step 3: Write the name of the cation first, followed by the anion. Na+  sodium ion PO43-  phosphate ion sodium phosphate

Writing the formula for polyatomic ions is the same as writing the formula for ionic compounds. You will use the cross-over method and your polyatomic table.

Na+ SO42- Na2SO4 Example: Write the name of sodium sulfate.
Step 1: Identify the cation. Step 2: Identify the anion. Step 3: Cross-over the charges to write the formula. sodium  Na+ sulfate  SO42- Na+ SO42- Na2SO4

Write the name of ammonium thiocyanate.
Example: Write the name of ammonium thiocyanate. NH4SCn See white board

Stock Naming System Most of the transition metals have more than one possible ion charge. They are often referred to as being multivalent. For example, Ion Possible Ion Charges Copper 1+, 2+ Iron 2+, 3+ Cobalt Chromium Lead 2+, 4+ Tin During the 18th century, German chemist Alfred Stock developed a new system called the Stock system. I n the Stock system, the "ous" and "ic" are replaced by Roman numerals representing the valence number of the atom. I n this system, ferrous becomes Fe(II) and ferric becomes Fe(III). A table showing ions with more than one valence is shown below. The Stock system solved the "ous"-"ic" problem, as you can see in the examples below. Fe(Cl)2 Iron (II) chloride Fe(Cl)3 Iron (III) chloride CuCl Copper (I) chloride Cu(Cl)2 Copper (II) chloride There is no confusion over what valence number a name is describing since the valence number of the metal ion is included in the name. Note the valence number of the metal can be determined from the number of non-metal ions. In the example, Fe(Cl)2 Iron II chloride there are two chlorines, indicating the valence number of the metal Fe.

In 1919, Alfred Stock (1876 – 1946), a German chemist, suggested using numbers to indicate the charge of the ions. Prior to this the ions were given different names based upon their charge. The Cu+ ion was called cuprous and the Cu2+ ion was called cupric. However, the Fe2+ ion was ferrous and the Fe3+ ion was ferric. Since the charges were not always the same, the "–ic" and "–ous" suffixes caused some confusion. Today, the Stock naming system uses Roman numerals following the metal ion's name to indicate the ion's charge.

Example: Copper (I) = Cu+ Copper (II) = Cu2+ Iron (II) = Fe2+ Iron (III) = Fe3+ cuprous cupric ferrous ferric

As a general rule, all metals are multivalent (have more than one ion charge) except group one and two metals, silver, cadmium, zinc, and aluminum. Unless the metal is one of these use the Roman numeral.

Example: Write the formula for iron (III) chloride Step 1: Write out the ions. Step 2: Cross-over the charges. Fe3+ and Cl- FeCl3

Example: Write the formula for lead (IV) sulfide. Step 1: Write the ions. Step 2: Cross-over the charges. Step 3: Reduce the subscripts. Pb4+ and S2- Pb2S4 Pb2S4 ÷ 2  PbS2 *divide by GCF*

Naming Multivalent Compounds
We name in a very similar manner as those ions with a single ion charge, except we must determine the charge on the metal ion. See overhead for examples

Naming Covalent Compounds
Non-metals tend to combine chemically by sharing electron pairs. These bonds are known as covalent bonds. Neutral compounds made of atoms joined covalently are called molecular or covalent compounds. Covalent compounds are also named i n the same way as ionic compounds (e.g., hydrogen fluoride is a predictable name for the compound formed from hydrogen and fluorine). The first nonmetal element is named and the second non-metal element is named with the suffix "ide" added. Some compounds also use common names. These names (e.g., water for H20 ) have become such a part of life that changing to dihydrogen oxide would be difficult. Other common names include ammonia for NH3 , and methane for CH4 . Descriptive prefixes are used to identify molecules according to IUPAC rules. Sometimes elements combine in more than one way, so more descriptive names are used to help identify the compound (e.g., carbon monoxide and carbon dioxide). In carbon monoxide, carbon combines with one oxygen; in carbon dioxide, carbon combines with two oxygens.

We name covalent compounds differently than ionic compounds
We name covalent compounds differently than ionic compounds. We must indicate the number of each element by adding a prefix in front of the element's name. Subscript Prefix one mono two di three tri four tetra five penta six hexa seven hepta eight octa nine nona ten deca

H2O = water NH3 = ammonia CH4 = methane
There are three exceptions to the naming rules. Here the common names for the compounds are used: H2O = water NH3 = ammonia CH4 = methane

Example: Write the name for CO2. Step 1: Name the first atom with prefixes. Step 2: Name the second element using prefixes and end in "-ide". Step 3: Write the name of the compound writing the substance found more to the left on the periodic table first. There is only 1 carbon. We omit “mono” for the first element. carbon There are 2 oxygens, so we use the di prefix dioxide carbon dioxide

Example: Write the name for N2O4. dinitrogen tetraoxide Example: Write the name for SF6. sulfur hexafluoride

Writing Covalent Formulas
Writing formulas for covalent compounds involves the following rules: Write the symbol for the first element followed by the subscript indicated by the prefix. Write the symbol of the second element followed by the subscript indicated by its prefix. Writing formulas for covalent compounds follows the same pattern as ionic compounds. When both covalent and ionic bonds are formed, they use electrons in the outer shells (valence electrons). Covalent bonds are formed when two non-metallic elements combine to form a compound. 1. Write both symbols, left-most element in the periodic table first, with valence numbers as shown. 2. Exchange valence numbers and place them as subscripts, as shown in the diagram below. 3. Reduce the formula 4. The subscript " 1 " is dropped. DO NOT REDUCE THE SUBSCRIPTS!!!

1 as a subscript is not needed
Example: Write the formula for dinitrogen monoxide. Step 1: Write the symbol and subscript for the first element. Step 2: Write the symbol and subscript for the second element. Step 3: Combine dinitrogen  N2 1 as a subscript is not needed monoxide  O N2O

SCl6 CCl4 Example: Write the formula for sulphur hexachloride.
Write the formula for carbon tetrachloride. CCl4 Naming Covalent Compounds 1 How to Name Ionic & Covalent Compounds Naming Covalent Compounds 2

Balancing Equations The Law of Conservation of Mass tells us that every chemical equation must have equal numbers of atoms of each element on each side of the equation. This means atoms cannot be created or destroyed in a chemical reaction.

Top-The seesaw is not balanced
Top-The seesaw is not balanced. There are seven atoms on the left and only six on the right. Bottom-The number of atoms is now in balance. The mass of the reactants also equals the mass of the products.

Four atoms of hydrogen and two atoms of oxygen.

As you balance equations, there are several rules to remember
You cannot change the formula of any reactant or product to change the numbers of atoms. You can only change the coefficients in front of the reactants and products. As you balance equations, there are several rules to remember: 1. You cannot change the formula of any reactant or product to change the numbers of atoms. 2. You c an only change the coefficients in front of the reactants and products. Coefficients can be placed only in front of the formula, not somewhere inside it ; that is, you can write 2H20, but not H220. b) Coefficients apply to the whole molecule; that is, 2H20 means there are two molecules of H 2 0 , which, in turn, means there are four atoms of hydrogen and two atoms of oxygen in those two molecules. If no coefficient is shown in front of a molecule, it means the molecule has a coefficient of 1. c) Note the difference between a coefficient and a subscript. A coefficient tells us the number of molecules in an equation and can be changed to balance an equation. A subscript tells us the number of atoms in a molecule and cannot be changed. The formula for water is H20 . The subscript 2 following the H means there are two hydrogen atoms combining with one oxygen atom to form one water molecule and it cannot be changed.

Note the difference between a coefficient and a subscript.
Coefficients can be placed only in front of the formula, not somewhere inside it Coefficients apply to the whole molecule. If no coefficient is shown in front of a molecule, it means the molecule has a coefficient of 1. Note the difference between a coefficient and a subscript. A coefficient tells us the number of molecules in an equation and can be changed to balance an equation. A subscript tells us the number of atoms in a molecule and cannot be changed.

There are some rules that you can use when you are given an unbalanced
H2(g) + O2(g)  H2O(l)

The number of atoms on each side of the equation are not the same.
Determine the number of atoms for each element in the molecules for reactants and products. The number of atoms on each side of the equation are not the same.

If the numbers of atoms on both sides of the equation are equal at this point, the equation is already balanced and you are finished. In this example, they are not the same and you go to the next step.

Choose the substance that has the most influence on the equation and insert coefficients in the formulas as needed.

Inspect the equation and recalculate the numbers of atoms on both sides of the equation. If they are equal, the equation is balanced. If they are not equal, change the coefficients until the equation is balanced. A Beginner’s Guide to Balancing Equations How to Write Chemical Equations How to Balance Equations How to Write, Balance & Classify a Chemical Reaction

Chemical Reactions In chemistry, a reaction happens when two or more molecules interact and the molecules change. Chemical reactions are an important part of our daily lives. When you start a car engine, gasoline burns, producing the energy needed to drive down the street. When you (or the baker) make a cake, ingredients are mixed, a chemical reaction takes place, and carbon dioxide makes bubbles in the cake. If it were not for chemical reactions, all cakes would be rather heavy! What goes in (reactants) always weighs the same as what comes out (products).

Single Displacement Reactions
In single displacement reactions, one element replaces another element in a compound. There are two possible reactions: One positive ion replaces another Zn + HCI  ZnCl2 + H2 b) One negative ion replaces another Cl2 + 2NaBr  2NaCl + Br2 Single Replacement Reactions In single replacement reactions, one element replaces another element i n a compound. There are two possible reactions. One positive ion replaces another. Zn + HCI  ZnCl2 + H2 Note that the zinc and hydrogen are both positive ions when they form molecules. Note that hydrogen is diatomic. It is a good idea to learn the diatomic elements in order to know that they form a diatomic molecule in some chemical reactions. b) One negative ion replaces another. Cl2 + 2NaBr  2NaCl + Br2 Single replacement reactions can be illustrated using the following general equation. element + compound  element + compound A + BC  B + AC Some practical examples of a single replacement reaction include: • placing aluminum foil in a solution of iron (III) nitrate • placing a copper wire in a silver nitrate solution An element is reacted with a compound. One element may replace another element to form a new compound. A + BC  B + AC

Synthesis Reactions A + B  AB
Synthesis reactions involve joining atoms to make a molecule, or joining elements to form a compound. Two or more simple elements or compounds combine to form a more complex compound. 2Mg + O2  2MgO Synthesis Reactions Two or more simple elements or compounds combine to form a more complex compound in synthesis reactions. A sample reaction is shown below. 2Mg + 02  2MgO A practical illustration of a synthesis reaction is to place powdered zinc and powdered sulfur in a fume chamber and heat with a Bunsen burner. A synthesis reaction can be illustrated using the general equation shown below. two or more elements or compounds  compound A + B  AB Involves joining atoms to make a molecule, or joining elements to form a compound. A + B  AB

Decomposition Reactions
During decomposition, one compound splits apart into two (or more) pieces. sodium chloride  sodium + chlorine 2HgO 2Hg + 02 Decomposition Reactions During decomposition, one compound splits apart into two (or more pieces). These pieces can be elements or simpler compounds. A sample reaction is shown below. 2HgO 2Hg + 02 A decomposition reaction can be illustrated using the general equation shown below. compound  two or more elements or compounds AB  A + B Reverse of synthesis-molecules/compounds breaking apart. AB  A + B

Double Displacement Reactions
In double displacement, elements in different compounds replace each other (split up and exchange partners). Double Replacement Reactions In a double replacement reaction, the two molecules split up and exchange ion partners. A sample reaction is shown below. FeS + 2HC1 FeCl2 + H2S A practical illustration of a synthesis reaction is to add potassium iodide solution to lead(II) nitrate solution. A double replacement reaction can be illustrated using the general equation shown below. compound + compound  compound + compound AC + BD  AD + BC In double displacement, elements in different compounds replace each other. AC + BD  AD + BC

Combustion Reactions Hydrocarbon + O2  CO2 + H20
Combustion reactions involve the burning of a chemical substance in oxygen. Metals burn in oxygen to produce a metal oxide. Magnesium + oxygen  magnesium oxide Nonmetals burn in oxygen to produce a dioxide. Carbon + oxygen  carbon dioxide Combustion Reactions Combustion, at its most general, can mean the reaction of oxygen gas (02) with another material. We understand combustion to mean the reaction of oxygen with a compound containing carbon and hydrogen. A common synonym for combustion is "burn." An example combustion reaction is shown below. To test for carbon dioxide, place a burning splint in a container of the gas. The splint will go out since carbon dioxide will not support combustion. CH4 is a hydrocarbon called methane. The common name of the hydrocarbon is not important here. What does matter is that all hydrocarbons combine w i t h oxygen when they burn to produce water and carbon dioxide. A practical illustration of a combustion reaction occurs when we light a candle or eat some food. A combustion reaction can be illustrated using the general equation shown below. hydrocarbon + oxygen  carbon dioxide + water CH C02 + H20 C02 + H2 Involves the burning of a chemical substance in oxygen. Metals burn in oxygen to produce a metal oxide. Non metals burn in oxygen to produce a dioxide. Hydrocarbon + O2  CO2 + H20

Bill Nye Chemical Reactions
Exothermic and Endothermic Reactions 7 Chemical Reactions that will Facinate

Common Acids and Bases Acids are used in many industrial processes. Production of paper, steel, and many other products requires the use of acids. Sulphuric and hydrochloric acids are commonly used in such industrial processes. Bases are important household chemicals used for cleaning and disinfecting. Bases are also used in hairdressing when a "permanent" change is needed to change straight hair to curly hair. You have already experienced the sour taste of acids. We like sour tastes. Lemon juice in water with a little honey has been used as a refreshing drink for centuries. We like carbonated beverages that make use of carbonic acid. Pickles need vinegar (acetic acid) to prevent them from spoiling. As you can see, acids are an important part of our daily lives. Acids: -low pH <7 -ability to lose hydrogen ions Bases: -high pH >14 -most common bases are alkalis (groups 1+2)

The names and formulas of some common acids and bases are shown below.
hydrochloric acid HCl sodium hydroxide NaOH sulphuric acid H2SO4 calcium hydroxide Ca(OH)2 nitric acid HNO3 ammonium hydroxide NH4OH Acids are used in many industrial processes. Production of paper, steel, and many other products requires the use of acids. Sulfuric and hydrochloric acids are commonly used i n such industrial processes. Bases are important household chemicals used for cleaning and disinfecting. Bases are also used in hairdressing when a "permanent" change is needed to change straight hair to curly hair. Notice that acids tend to have H in their formula and bases often have OH (hydroxide) as the second component in their formula.

Characteristics of Acids and Bases
Taste Feel Reaction to Litmus Paper Other Properties sour bitter burns slippery turns litmus paper red turns litmus paper blue neutralizes basic solutions corrosive to metal conduct electricity neutralizes acidic solutions conduct electricity

Reactions Examples Acid + bases  water + salt Acid + metals  H2 gas
Acid + carbonates  CO2 + H2O + salt Bases + acids  water salt NaOH – sodium hydroxide Ca(OH)2 – limewater Mg(OH)2 – milk of magnesia bleaches soaps toothpaste HCl – gastric juice citric acid –lemons & oranges acetic acid – vinegar ascorbic - vitamin C acid acetylsalicylic - Aspirin acid

A substance that changes colour when added to an acid or base is called an indicator.
Litmus red blue Phenolphthalein colourless pink Methyl Orange orange yellow Bromothymol Blue light yellow

acid + base  salt + water
Neutralization An acid and a base, when combined, will neutralize each other. Acids will lose their acid properties and bases will lose their base properties. When an acid reacts with a base, a salt and water are produced. acid + base  salt + water

HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
A salt is a compound composed of the negative ion of an acid and the positive ion of the base. The water is formed when the hydrogen ion (H+) of the acid combines with the hydroxide ion (OH-) of the base. acid base salt water HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

Measuring the Strength of acids and Bases
The strength of an acid or a base is measured using a scale called pH. pH 0-6.9 acidic Measures the strength of acid where a smaller number equals stronger acid and a larger number equals weaker acid pH 7 (water) neutral Neutral - neither acid nor base pH basic (alkaline) Measures the strength of base where smaller number equals weaker base and larger number equals stronger base

strong weak weak strong

Question: Use the pH values described below to determine whether the substance is neutral, a weak or strong acid, or a weak or strong base. pH = 12 pH = 7 pH = 6 pH = 9 pH = 2 strong base neutral weak acid weak base strong acid Acids, Bases and pH

UNIT 2 Chemistry in Action.

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