CHEMISTRYUNIT 7 REVIEW PACKET

MESA Charter High School

NAME _________________________Period _____________

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Learning Goal 7.1 – I can identify and describe the formation and behavior of ionic bonds.Review Notes:

• Electronegativity - a measure of the tendency of an atom to attract a bonding pair of electrons• Ionization Energy - the minimum amount of energy needed to remove an electron• EXOTHERMIC = when energy is RELEASED as a product • ENDOTHERMIC = when energy is ABSORBED or needed as an ingredient to fuel the process • Bonding creates STABILITY • Bond formation (spontaneous) = energy is released• Breaking Bonds (NOT spontaneous) = Energy is ABSORBED (consumed)

• Electrostatic attraction – the attraction between two opposite (+ and -) charges. This is what holds two ions together in an ionic bond.

• Ionic Compounds – cations (with a + charge) and anions (with a - charge) attract each other, creating an ionic bond.

• Ionic bonds are metal + non-metal

• Properties of Ionic Compounds• Exist as solids, and in aqueous solutions• High m.p./b.p (melting point/ boiling point) due to strong bonds • Electrolyte (conducts electricity when dissolved in an aqueous solution) • Hard (not brittle)• Sold ionic compounds are usually arranged in regular (neat, uniform) three dimensional

crystal lattices (networks).

Practice Questions

1. The element with the highest electronegativity on the periodic table is __________.a. Elements with high electronegativity have a high/low attraction for electrons.b. Elements with a high electronegativity are more likely to gain/lose electrons. c. Elements with high electronegativity are more likely to become cations/anions.

2. The element with the highest ionization energy on the periodic table is __________.a. Elements with a low ionization energy are more likely to gain/lose electrons. b. Elements with low ionization energy are more likely to become cations/anions.c. Elements with low ionization energy are more likely to become cations/anions.

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Learning Goal 7.2 – I can identify and describe the formation and behavior of covalent bonds and multiple covalent bonds.Review Notes:

• A Covalent Bond is a chemical bond that involves the sharing of electrons between atoms.• The two atoms are non-metal + non-metal• Covalent Bonds can also be called molecular bonds• Covalent Compound = molecular compound• Sometimes two elements can form more than one covalent bond

• Single bonds = 2 electrons shared• Double bonds = 4 electrons shared• Triple bonds = 6 electrons shared

• Properties of Covalent Compounds • Gases, liquids, or soft solids (made of molecules) there are exceptions!• Low m.p./b.p (melting point/ boiling point) because of low forces between atoms• Poor electrical conductors in all phases• Non-electrolytes (do not conduct electricity in H2O)• Brittle• Some are polar and partially soluble in water• Many soluble in nonpolar liquids but not in water

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Learning Goal 7.3 – I can identify and describe the formation and behavior of metallic bonds. Review Notes:

• Metallic bonds• 2 or more metals• Atoms of metals are tightly packed together in a giant lattice similar to the lattice in ionic

compounds.• The outer electrons separate from their atoms and become delocalized, creating a ‘sea

of electrons’. The atoms become positive and are attracted to these negative electrons.

Properties of Metalso High melting and boiling points (bonds are super strong)o Excellent conductors of electricity (solid and liquid) à more electrons, more

conductivityo Excellent Conductors of heat (solid and liquid) à increased vibrations spread through

sea of electronso Good Malleability and ductility à atoms in a metal can roll over each other while still

maintaining their bonds, especially if heated.1

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3. Which one of the following properties is not characteristic of typical metals?

a. Moderately high melting pointb. High boiling pointc. Brittled. Good electrical conductor when

a solid

4. Of the three major types of bonding (ionic, covalent, metallic) which ones are based on electrostatic attractions between opposite charges?

a. Only ionic bondingb. Only ionic and covalent bondingc. Only ionic and metallic bondingd. Ionic, covalent and metallic

bonding

5. True or false: metals give up electrons easilya. Trueb. False

6. Why are metals described as having a “sea of electrons”?

a. Electrons are wetb. Electrons move in wavesc. Electrons move freely between atomsd. Electrons are fixed to metal atoms

7. When 2 or more metals are combined they form an…

a. Brassb. Alloyc. Bronzed. Covalent bond

8. Which of the following is a correct description of a metallic property?

a. Metals are brittleb. Metals are malleablec. Metals have a dull appearanced. Metals gain electrons to form anions

9. Which of the following is not a property of metals?

a. Shiny lusterb. Brittle/shatters easilyc. Conducts electricityd. Malleable

10. Which of the following is a mixture of elements that have metallic properties

a. An alloyb. A suspensionc. A gasd. A pure metal

11. What does the “sea of electrons” contain?a. All the electrons in that metalb. The electrons in the outer shell of that

metalc. The electrons in the inner shell of that

metal

12. Why are metals able to conduct electricity?a. The positive metal ions pass charges to

each otherb. Electrons pass charges through the

positive metal ionsc. The sea of electrons helps pass changes

through the metal

13. Metallic bonding occurs between atoms ofa. Sulfurb. Copperc. Fluorined. Carbon

14. The high electrical conductivity of metals is primarily due to

a. High ionization energiesb. Full valence shellsc. Mobile electronsd. High electronegativity

15. A substance that conducts electricity as a solid and when melted into a liquid is most likely classified as

a. An ionic compoundb. A molecular compoundc. A metald. A nonmetal

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Learning Goal 7.4 – I can draw Lewis dot structures for molecules and compounds.

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Learning Goal 7.5 – I can explain polar bonds and polar molecules in terms of electronegativity; determine the degree of polarity of a bond based on electronegativity, and describe the shape of a molecule based on its polarity.Review Notes:

Polar Covalent Bondso The bond is covalent – the electrons are still shared between the 2 atoms. But the

electrons are NOT shared equally.o One atom has a much higher electronegativity than the othero Neither atom in a polar bond is entirely positive or entirely negative – instead we say

partial positive (δ+) and partial negative (δ-). Differences in electronegativity predict the type of bond formed

o Greater than 2.0 bond is ionico Between 2.0 and 0.4, bond is polaro Less that 0.4, bond is non-polar

Like Dissolves Likeo Polar compounds only dissolve in polar solutionso Non-polar compounds only dissolve in non-polar solutions

Drawing Polar Compoundso Draw the Lewis Dot Structure for the compoundo Mark partial positive/negative areas (using Table S)o If the bonds are non-polar, no δ+/- neededo Use “polar bond arrows” <+ to point at the partial negative end of the bond.

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Draw the Lewis Dot Structures for the following (possibly polar) compounds. Label partial positive and partial negative atoms. Include polarity arrows next to each polar bond. Use symmetry to determine if the molecule is polar or non-polar.

CH4 PF3

1) Does this molecule have polar bonds?

2) Is this molecule symmetrical in X and Y?

3) Is this molecule polar?

1) Does this molecule have polar bonds?

2) Is this molecule symmetrical in X and Y?

3) Is this molecule polar?

CH3F CH2Br2

1) Does this molecule have polar bonds?

2) Is this molecule symmetrical in X and Y?

3) Is this molecule polar?

1) Does this molecule have polar bonds?

2) Is this molecule symmetrical in X and Y?

3) Is this molecule polar?

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H2S CBr4

4) Does this molecule have polar bonds?

5) Is this molecule symmetrical in X and Y?

6) Is this molecule polar?

4) Does this molecule have polar bonds?

5) Is this molecule symmetrical in X and Y?

6) Is this molecule polar?

PCl3 H2O

4) Does this molecule have polar bonds?

5) Is this molecule symmetrical in X and Y?

6) Is this molecule polar?

4) Does this molecule have polar bonds?

5) Is this molecule symmetrical in X and Y?

6) Is this molecule polar?

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Learning Goal 7.6 – I can explain how unequal distribution of charge results in intermolecular forces, including van der Waals forces and hydrogen bonds. Review Notes:

The forces holding molecules together. Intermolecular forces are NOT bonds. They are much weaker than all types of bonds. As IMFs get larger (stronger), melting point and boiling points increase Van der Waals forces - a name for all of the different attractive or repulsive forces between

molecules Ion-Dipole Forces - Attractive force resulting from the electrostatic attraction between an ion

and a neutral molecule that has a dipole. Dipole-Dipole Forces - Occur when two polar molecules approach each other δ+ of one

molecule attracted to δ- of the othero Attractive force gets larger as molar mass increases

London Dispersion Forces - Weakest intermolecular force. Only important for non-polar molecules

o Electron-electron repulsion creates temporary dipoles that attract each othero More electrons = greater London dispersion forces

Hydrogen Bonds - (not actually bonds) The electrostatic attraction between polar molecules. o Occurs when an H atom bound to a highly electronegative atom (S, N, O, or F) is also

attracted to another highly electronegative atomo H-bond donor: the atom to which the the hydrogen atom is covalently boundo H-bond acceptor: the neighboring electronegative ion or molecule

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Learning Goal 7.7 – I can describe the differences in structure, formation and properties of ionic and molecular compounds.

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Properties IONIC COVALENT METALLIC

Action of electrons Molecular Network

Type of element

Bond strength

Melting point / boiling pointHard vs. Brittle

Soluble in H2O

Electrolytes?

Conduct electricity as a solid?Conduct electricity as a liquid?Conducts heat?

Multiple compounds are held together by

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Learning Goal 7.8 – I can explain the properties of giant covalent compounds in terms of their structures. Review Notes:

• Ionic solids are formed by linking individual compounds with intermolecular forces. Thus the individual compounds are not actually bonded together.

• Metallic and Covalent solids are formed solids by joining multiple compounds together with additional covalent bonds – thus there is continuous bonding throughout the solid structure.

• Network Solids - Large 3D networks of covalent compounds connected by covalent bonds. • Allotropes = “other form” - 2 or more compounds with the same composition, but different

arrangements of bonds. • Properties of Network Solids

• Strong covalent bonding• Extended, highly regular network of atoms• Do NOT conduct electricity à all valence electrons are trapped in covalent bonds• High m. p. and b. p.à very high amount of energy required to break such a huge number of

covalent bonds at once. • Insoluble in water à bonds are all non-polar

• Graphite• Each carbon is bonded to 3 other carbon atoms. The 4th valence electron becomes

delocalized, and is free to move à thus graphite DOES conduct electricity!• Network of carbon atoms are arranged in sheets (layers). The sheets are held together by

IMFs.• Layers can be scraped off à used in pencils.• Melting point = 5800K ( or at high pressure)• Hardness = 1-2

• Diamond• 4 single covalent bonds to each carbon• Hardest known natural substance• Melting point = 4073 K• Hardness = 10• Does not conduct electricity (no free electrons!)

• Quartz - Silicon Dioxide• Extended network of silicon-oxygen bonding (SiO2)• Melting point = 1600°C• Hardness = 7• Amethyst is quartz with the addition of Iron3+ cations • Does not conduct electricity (no free electrons!)

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