Additional Chapter 5 information Write on a piece of paper for notes: Ion – an atom or group of bonded atoms that has a positive or negative charge. (not neutral anymore) Ionization – Any process that results in the formation of an ion A + energy  A+ + e- (atom loses an e- and becomes more positive) A + e-  A- + energy (atom gains an e- and becomes more negative)

Ionic Radii Cation – A positive Ion. Formed by the loss of one or more electrons. Always leads to a decrease in atomic radius. KITTY …. KITTY …. KITTY  Anion – A negative Ion. Formed by adding one or more electrons. Always leads to an increase in atomic radius. https://www.youtube.com/watch?v=900dXBWgx3Y

Comparing Cations and Anions Section 3 Electron Configuration and Periodic Properties Chapter 5 Comparing Cations and Anions Click below to watch the Visual Concept.

Valence Electrons – electrons available to be lost, gained, or shared in the formation of chemical compounds. these electrons are in the highest energy level and are often located in incompletely filled main-energy levels. For Main-Group elements, the valence e- are the electrons in the outermost s and p sublevels *Table 4 p. 160

Section 3 Electron Configuration and Periodic Properties Chapter 5 Valence Electrons Click below to watch the Visual Concept. Visual Concept

Chap 6: Introduction to Chemical Bonding Learning Targets: 1) Define chemical bond. 2) Explain why most atoms form chemical bonds. 3) Describe ionic and covalent bonding. 4) Explain why most chemical bonding is neither purely ionic nor purely covalent. 6) Classify bonding type according to electronegativity differences.

Introduction to Chemical Bonding Chemical Bond – is a mutual electrical attraction between the nucleus and valence electrons of different atoms that binds the atoms together. By themselves, most atoms have a high potential energy. Nature favors arrangements in which potential energy is minimized. By bonding with each other, atoms decrease in potential energy, creating more stable arrangements of matter.

Chemical Bond Chapter 6 Section 1 Introduction to Chemical Bonding Click below to watch the Visual Concept. Visual Concept

Types of Chemical Bonding Ionic Bonding – chemical bonding that results from the electrical attraction between large numbers of cations and anions. (METALS) atoms completely give up electrons to other atoms (fig. 1 page 176) Covalent Bonding – results from the sharing of electron pairs between two atoms (fig. 1 p. 176) (NONMETALS) Nonpolar-covalent bond – bonding electrons are shared equally by bonded atoms, resulting in a balanced distribution of electrical charge. (Fig. 3 p. 177) Polar-covalent bond bonding electrons are NOT shared equally by bonded atoms resulting in an unbalanced distribution of electrical charge. The Polar bear DOES NOT share! (Fig. 3 p. 177)

Ionic Bonding Chapter 6 Section 1 Introduction to Chemical Bonding Click below to watch the Visual Concept. Visual Concept

Covalent Bonding Chapter 6 Section 1 Introduction to Chemical Bonding Click below to watch the Visual Concept. Visual Concept https://www.youtube.com/watch?v=VSc491HLzDo Comparing ionic and covalent bonds

Comparing Polar and Nonpolar Covalent Bonds Section 1 Introduction to Chemical Bonding Chapter 6 Comparing Polar and Nonpolar Covalent Bonds Click below to watch the Visual Concept. Visual Concept

- Bonding between atoms is never purely ionic and is rarely purely covalent. - Electronegativity – measure of an atom’s ability to attract electrons. (chap. 5 p. 161) - The degree to which bonding between atoms of two elements is ionic or covalent can be estimated by calculating the difference in the elements’ electronegativities. sample problem A p. 177 (3 slides)

- Fig. 2 p. 176 – shows Electronegativity from 0 – 0.3 = nonpolar-covalent bond (0 – 5% ionic) Electronegativity from 0.3 – 1.7 = polar –covalent bond (5 – 50% ionic) Electronegativity from 1.7 – 3.3 = Ionic bonds (50 – 100% ionic) - Bonding between two atoms of the same element is completely COVALENT (called a diatomic element) - Carbon is unique because it can form 4 covalent bonds and bond with other carbon atoms to form long chain molecules of different sizes and shapes.

Using Electronegativity Difference to Classify Bonding Section 1 Introduction to Chemical Bonding Chapter 6 Using Electronegativity Difference to Classify Bonding Click below to watch the Visual Concept. Visual Concept

Chemical Bonding, continued Section 1 Introduction to Chemical Bonding Chapter 6 Chemical Bonding, continued Sample Problem A Use electronegativity values listed in Figure 20 from the previous chapter in your book, on page 161, and Figure 2 in your book, on page 176, to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative?

Chemical Bonding, continued Section 1 Introduction to Chemical Bonding Chapter 6 Chemical Bonding, continued Sample Problem A Solution The electronegativity of sulfur is 2.5. The electronegativities of hydrogen, cesium, and chlorine are 2.1, 0.7, and 3.0, respectively. In each pair, the atom with the larger electronegativity will be the more-negative atom. Bonding between Electroneg. More-neg- sulfur and difference Bond type ative atom hydrogen 2.5 – 2.1 = 0.4 polar-covalent sulfur cesium 2.5 – 0.7 = 1.8 ionic sulfur chlorine 3.0 – 2.5 = 0.5 polar-covalent chlorine

Section 2: Covalent Bonding and Molecular Compounds Learning Targets: Define molecule and molecular formula. 2) Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy. 3) State the octet rule. 4) Construct basic Lewis structures. 5) Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both. 6) Understand that molecules, with covalent bonds are always NONMETALS.

Covalent Bonding and Molecular Compounds Molecule – a neutral group of atoms that are held together by covalent bond nonmetals could be two or more of the same atom (O2) or two or more different atoms (C6H12O6) (nonmetals) Diatomic Molecule – molecule containing only two atoms (O2 ; CO) A molecule is the smallest possible amount of a compound

Visual Concepts Chapter 6 Molecule

Comparing Monatomic, Diatomic, and Polyatomic Molecules Visual Concepts Chapter 6 Comparing Monatomic, Diatomic, and Polyatomic Molecules

Molecular Compound – a chemical compound whose simplest units are molecules Chemical Formula – indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts Molecular Formula – shows the types and numbers of atoms combined in a single molecule of a molecular compound. (H2O) (nonmetals)

Chemical Formula Chapter 6 Section 2 Covalent Bonding and Molecular Compounds Chapter 6 Chemical Formula Click below to watch the Visual Concept. Visual Concept

Formation of a Covalent Bond Fig. 5 p. 179 A) Separated Hydrogen atoms each have high Potential Energy (PE) B) PE decreases as atoms are drawn together by attractive forces.(Protons are attracted to electrons) There is also Repulsive forces (P vs P and e- vs e-) C) PE is at a minimum when the attractive forces are balanced with the repulsive forces. D) PE increases when repulsion between like charges is greater than the attraction between opposite charges.

Formation of a Covalent Bond Section 2 Covalent Bonding and Molecular Compounds Chapter 6 Formation of a Covalent Bond

Formation of a Covalent Bond Section 2 Covalent Bonding and Molecular Compounds Chapter 6 Formation of a Covalent Bond

Characteristics of the Covalent Bond Bond Length – the distance between 2 bonded atoms at their minimum PE Bond Energy – the energy required to break a chemical bond and form neutral isolated atoms. Measured in Kilojoules per mole. Bond Length is inversely proportional to Bond Energy, so as Bond length gets shorter the bond energy gets larger

Bond Energies and Bond Lengths for Single Bonds Section 2 Covalent Bonding and Molecular Compounds Chapter 6 Bond Energies and Bond Lengths for Single Bonds

Section 2 Covalent Bonding and Molecular Compounds Chapter 6 Bond Length Click below to watch the Visual Concept. Visual Concept

Section 2 Covalent Bonding and Molecular Compounds Chapter 6 Bond Energy Click below to watch the Visual Concept. Visual Concept

Noble Gases – Outer s and p orbitals are completely The Octet Rule – Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an Octet of e- in its highest occupied energy level. Noble Gases – Outer s and p orbitals are completely filled with 8 electrons making them very stable. (except He – has full 1s) * Other main group atoms can effectively fill their outermost s and p orbitals with e- by sharing e- through covalent bonding by following the Octet Rule. *Fig. 9 p. 183 There are some exceptions to the Octet Rule (p. 183)

Section 2 Covalent Bonding and Molecular Compounds Chapter 6 The Octet Rule Click below to watch the Visual Concept. Visual Concept https://www.youtube.com/watch?v=ZxWmyZmwXtA

- For the main group elements – don’t include the d block elements Electron – Dot Notation – an electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol. - For the main group elements – don’t include the d block elements - Must be in this order: 3 7 X1 2 X X X X5 6X X X 4 8 Let’s Practice! P Si S Cl Xe

Section 2 Covalent Bonding and Molecular Compounds Chapter 6

Lewis Structures - formulas in which atomic symbols represent nuclei and inner-shell e- (F) - Dot-pairs or Dashes between two atomic symbols represent e- pairs in covalent bonds H - H or F : F - Dots adjacent to only one atomic symbol represent unshared e- :F : F: or :F - F: - Structual Formula – indicates the kind, number, arrangement, and bonds but not the unshared pairs of atoms in a molecule. F - F H-Cl - Single Bond – a covalent bond produced by the sharing of one pair (two) electrons between two atoms - Follow the steps in sample problem C p. 185 on how to write a Lewis Structure.

Lewis Structures Chapter 6 Section 2 Covalent Bonding and Molecular Compounds Chapter 6 Lewis Structures Click below to watch the Visual Concept. Visual Concept

Double Bond – covalent bond produced by the Multiple Covalent Bonds - some elements, epecially Carbon, Nitrogen, and Oxygen can share more than one electron pair. Double Bond – covalent bond produced by the sharing of two pairs (4) electrons between two atoms Triple Bond – covalent bond produced by the sharing of three pairs (6) electrons between two atoms. **The need for a multiple bond becomes obvious if there are not enough valence e- to complete octets by adding unshared pairs. *Look at sample problem D p. 188 Resonance Structures – read p. 189

Comparing Single, Double, and Triple Bonds Section 2 Covalent Bonding and Molecular Compounds Chapter 6 Comparing Single, Double, and Triple Bonds Click below to watch the Visual Concept. Visual Concept

Section 6-3 Ionic Bonding and Ionic Compounds Objectives “I can” statements: 1) Compare a chemical formula for a molecular compound with one for an ionic compound.  2) Discuss the arrangements of ions in crystals.   3) Define lattice energy and explain its significance. 4) List and compare the distinctive properties of ionic and molecular compounds.

Section 6-3 Ionic Bonding and Ionic Compounds Ionic Compound – composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. Most Ionic compounds exist as crystalline solids. The chemical formula of an ionic compound represents the simplest ratio of the compound’s combined ions that gives electrical neutrality. Na+ Cl-  NaCl Molecular compounds are independent, neutral units with covalent bonds that can be isolated and examined Formula Unit – the simplest collection of atoms from which an ionic compound’s formula can be established. the ratio of ions in a formula unit depends on the charges of the ions combined. Ex: above with NaCl Ex: CaF2 (F- Ca2+)

Formation of Ionic Compounds Ca + F + F  Ca2+ + F- + F- = (CaF2) The Calcium atom has two valence electrons and each Fluorine atom has seven valence electrons. Atoms of Calcium and other alkaline earth metals easily lose two electrons to form cations.  - Atoms of Fluorine and other halogens easily gain one electron to form anions. The transfer of two electrons from Ca to each F atom transforms each atom into an ION with noble-gas configuration.

https://www.youtube.com/watch?v=zpaHPXVR8WU

Characteristics of Ionic Bonding In an Ionic Crystal, ions minimize their Potential Energy (PE) by combining in an orderly arrangement known as CRYSTAL LATTICE.

NaCl and CsCl Crystal Lattices Section 3 Ionic Bonding and Ionic Compounds Chapter 6 NaCl and CsCl Crystal Lattices

Characteristics of Ion Bonding in a Crystal Lattice Section 3 Ionic Bonding and Ionic Compounds Chapter 6 Characteristics of Ion Bonding in a Crystal Lattice Click below to watch the Visual Concept. Visual Concept

The distances between ions and their arrangement in a crystal represent a balance between attractive and repulsive forces NaCl – in the crystal structure each Na is surrounded by 6 Cl and each Cl is surrounded by 6 Na for a 1:1 ratio. Fig. 15 p. 192.

Lattice Energy – the energy released when one mole of an ionic crystalline compound is formed from gaseous ions. This determines the IONIC BOND STRENGTH

Lattice Energy Chapter 6 Section 3 Ionic Bonding and Ionic Compounds Click below to watch the Visual Concept. Visual Concept

Comparision of Ionic and Molecular Compounds The force that holds ions together in ionic compounds is a very strong ELECTROSTATIC attraction between positive and negative charges. The forces between covalent bonds in molecules is much weaker than the forces in ionic bonding. Ionic compounds generally have higher melting and boiling points than molecular compounds. Ionic compounds are hard but brittle. A slight shift of one row of ions relative to another causes a large buildup of repulsive forces. If one layer is moved, the repulsive forces make the layers part completely and break (brittle). (Fig. 17 p. 193)

Melting and Boiling Points of Compounds Section 3 Ionic Bonding and Ionic Compounds Chapter 6 Melting and Boiling Points of Compounds

Ionic Vs. Covalent Bonding Section 3 Ionic Bonding and Ionic Compounds Chapter 6 Ionic Vs. Covalent Bonding

Many Ionic compounds are easily dissolved in water and become good electrical conductors. Don’t conduct electricity in the crystal state.

Comparing Ionic and Molecular Compounds Section 3 Ionic Bonding and Ionic Compounds Chapter 6 Comparing Ionic and Molecular Compounds Click below to watch the Visual Concept. Visual Concept

Polyatomic Ions – a charged group of covalently bonded atoms have both molecular (covalent) and ionic characteristics the charge of a polyatomic ion results from an excess of electrons (negative charge) or a shortage of electrons (positive charge). examples on p. 194

Polyatomic Ions, continued Section 3 Ionic Bonding and Ionic Compounds Chapter 6 Polyatomic Ions, continued Some examples of Lewis structures of polyatomic ions are shown below.

Section 6-4 Metallic Bonding Objectives or “I can” Statements: 1) Describe the electron-sea model of metallic bonding, and explain why metals are good electrical conductors. 2) Explain why metal surfaces are shiny. 3) Explain why metals are malleable and ductile but ionic-crystalline compounds are not.

Section 6-4 Metallic Bonding Metallic Bonding – The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons. Vacant orbitals in the atoms’ outer energy levels overlap. This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal. The electrons are delocalized, which means that they do not belong to any one atom but move freely about the metal’s network of empty atomic orbitals. These mobile electrons form a “sea of electrons” around the metal atoms which are packed together in a crystal lattice.

Metallic Bonding Chapter 6 Section 4 Metallic Bonding Click below to watch the Visual Concept. Visual Concept

Metallic Properties High electrical and thermal conductivity results in metals due to the freedom of motion of electrons in a network. Shiny appearance is the result of metals being able to absorb a wide range of light frequencies. Electrons are excited and fall back to lower E levels causing the shiny appearance. Metallic bonding is the same in all directions through out the solid. One plane of atoms in a metal can slide past another without any resistance or breaking any bonds. Allows special properties: Malleability – hammered into thin sheets Ductility – drawn or pulled to make wire

Properties of Metals: Surface Appearance Section 4 Metallic Bonding Chapter 6 Properties of Metals: Surface Appearance Click below to watch the Visual Concept. Visual Concept

Properties of Metals: Malleability and Ductility Section 4 Metallic Bonding Chapter 6 Properties of Metals: Malleability and Ductility Click below to watch the Visual Concept. Visual Concept

Properties of Metals: Electrical and Thermal Conductivity Section 4 Metallic Bonding Chapter 6 Properties of Metals: Electrical and Thermal Conductivity Click below to watch the Visual Concept. Visual Concept

Section 6-5 Intermolecular Forces Objectives or “I can” statements: Describe dipole-dipole forces and hydrogen bonding, and their effects on properties such as boiling and melting points. 2) Distinguish between Intramolecular bonds that hold the inside of a molecule together and Intermolecular bonds that bond individual molecules together.

Section 6-5 Intermolecular Forces Intermolecular Forces – forces of attraction between molecules vary in strength but are generally weaker than the bonds that join the molecules (intramolecular forces). Add this the higher the boiling point - the stronger the forces between particles. Boiling points for ionic compounds and metals tend to be much higher than those for molecular substances: forces between molecules are weaker than those between metal atoms or ions. Strongest Intermolecular forces exist between Polar Molecules

Comparing Ionic and Molecular Substances Section 5 Molecular Geometry Chapter 6 Comparing Ionic and Molecular Substances

Section 6-5 Intermolecular Forces Dipole-Dipole Forces are the forces of attraction between polar molecules - A dipole is created by equal but opposite charges that are separated by a short distance. - The direction of a dipole is from the dipole’s positive pole to its negative pole. - The negative region in one polar molecule attracts the positive region in adjacent molecules. So the molecules all attract each other from opposite sides.

Dipole-Dipole Forces Chapter 6 Section 5 Molecular Geometry Click below to watch the Visual Concept. Visual Concept

Section 6-5 Intermolecular Forces Hydrogen Bonding – a strong type of Dipole-Dipole force. Water molecules have a high boiling point due to this strong bond. Hydrogen bonding occurs between hydrogen atoms and very electronegative atoms which creates a POLAR bond. These elements are Nitrogen, Oxygen, and Fluorine. (N, O, F)

Visual Concepts Chapter 6 Hydrogen Bonding

London Dispersion Forces Section 5 Molecular Geometry Chapter 6 Intermolecular Forces, continued London Dispersion Forces - In any atom or molecule—polar or nonpolar—the electrons are in continuous motion. - As a result, at any instant the electron distribution may be uneven. A momentary uneven charge can create a positive pole at one end of an atom of molecule and a negative pole at the other.

London Dispersion Force Section 5 Molecular Geometry Chapter 6 London Dispersion Force Click below to watch the Visual Concept. Visual Concept

Intermolecular Forces, continued Section 5 Molecular Geometry Chapter 6 Intermolecular Forces, continued London Dispersion Forces, continued - This temporary dipole can then induce a dipole in an adjacent atom or molecule. The two are held together for an instant by the weak attraction between temporary dipoles. - The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles are called London dispersion forces.