1. Periodic Trends Bellringer
1) Which has a larger size, a Mg atom or a Mg ion?
2) smaller ionization energy, K or Br?
3) smaller size, F atom or an I atom?
4) smaller electronegativity, O or Se?
5) less shielding, Ca or Sr?

2. Answers
1) Mg atom
2) K
3) F
4) Se
5) Ca

3. Ch. 7 Ionic and Metallic Bonding

4. Valence Electrons
Valence electrons = electrons in the highest occupied energy level of an element’s atoms
This # largely determines the chem. properties of the element
To find the # of valence e-’s in an atom of a representative element, look at its group #
Group 1 has 1, group 2 has 2, group 13 or 3A has 3, …
Electron dot structures (Lewis dot diagrams) = diagrams that show valence electrons as dots
*Draw Lewis dot structures*

5. Octet Rule
Octet rule = atoms gain or lose electrons to achieve a stable level of usually 8
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 to share electrons w/ another nonmetal to achieve a complete octet

6. Cations
Cations (+ charged atom) are formed when they lose one or more valence electrons (ionization) in order to become stable
Usually lose from 1 to 3 valence electrons tops
Can use electron configurations to illustrate the point but let’s use electron dot diagrams for ease
Exceptions are due to having to lose or gain too many valence electrons to achieve a noble gas state, so:
Some atoms attain a pseudo noble-gas electron configurations – Cu - copper, Ag - silver, Au - gold, Cd - cadmium
Two atoms are walking down the street.
Says one atom to the other, "Hey! I think I lost an electron!"
The other says, "Are you sure??"
"Yes, I'm positive!"

7. Anions
Anions (- charged atoms) are formed when they gain 1 or more valence electrons
Typically suffix is –ide
Halide ions (F-, Cl-, Br-, I-) are halogens that gain 1 e-
CP 7.1, PP 1-2 pg. 193, sect. assessment 7.1 pg

8. Common ions Cation Name Anion Name* H 1+ hydrogen H 1- hydride
Li 1+ lithium F 1- fluoride
Na 1+ sodium Cl 1- chloride
K 1+ potassium Br 1- bromide
Cs 1+ cesium I 1- iodide
Be 2+ beryllium O 2- oxide
Mg 2+ magnesium S 2- sulfide
Al 3+ aluminum
Ag 1+ silver
*The root is given in color.

10. Ionic Compounds
Anions and cations are held together by opposite charges.
Ionic compounds are called salts
Electrically neutral
Simplest ratio is called the formula unit
The bond is formed through the transfer of electrons (called an ionic bond)
Electrons are transferred to achieve noble gas configuration
Most are crystalline solids at room temp.
High melting points – large attractive forces result in a very stable structure
Good conductors of electricity when melted or dissolved in water
Overheard at the mall
Teen 1: Did you hear oxygen and magnesium got together?? Teen 2: OMg!
Your mama's so ugly
Your mama's so ugly...even Fluorine won't bind to her

11. Formulas
Chemical formula = shows the kinds and #s of atoms in the smallest representative unit of a substance
NaCl is the chem. formula for sodium chloride
Formula unit = the lowest whole-number ratio of ions in an ionic compound
MgCl2, NaCl, AlBr3
CP 7.2, PP pg. 196,
7.2 sect. assessment pg

12. Ch. 7 Bellringer
Write the Lewis electron-dot symbol for each of the following
A) sodium
B) fluorine
C) magnesium ion (Mg2+)
Write the chemical formula that results when the following pairs of ions combine to form an ionic bond
D) Mn4+ and O2-
E) Li1+ and Cl1-

13. Metallic Bonds and Properties
Metallic bonds = consist of the attraction of the free-floating valence electrons for the positively charged metal ions
How atoms are held together in the solid
Metals hold onto their valence electrons very weakly
Think of them as positive ions floating in a “sea of electrons”
Electrons are free to move through the solid
Metals conduct electricity
Malleable - hammered into shape (bend)
Ductile - drawn into wires
Electrons allow cations to slide by each other under pressure
Metals are crystalline structures and atoms are arranged in very compact and orderly patterns
*Cu vs. Cu compounds hammer demo – pg. 202*

14. Alloys
Alloys = mixtures composed of 2 or more elements, at least one of which is metal
Brass – copper and zinc
Sterling silver – silver and copper
Bronze – copper and tin
Steel – Fe, Cr, and others
Impt. b/c their properties are often superior to those of their component elements (usually cheaper as well)
*Making an alloy DEMO* - pg. 205

15. Ch. 8 Covalent Bonding

16. Molecules
Covalent bond = formed by sharing electrons b/w 2 or more atoms
Molecule = a neutral group of atoms joined by covalent bonds
Diatomic molecule = a molecule consisting of 2 of the same atoms
H2, N2, O2, F2, Cl2, Br2, I2 are the diatomic molecules in nature
Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds
Most molecular compounds are composed of atoms to 2 or more nonmetals
*Make a table comparing covalent bonding and ionic bonding*

18. Molecular Formula
Molecular formula = the chemical formula for a molecular compound
Shows how many atoms of each element a molecule contains
H2O, CO2, C2H6, O2
8.1 Sect. assessment, Pg

20. Covalent Bonding
Electron sharing usually occurs so that atoms attain the electron configuration of noble gases
Combos of atoms of the nonmetals and metalloids in 4A, 5A, 6A, and 7A are likely to form covalent bonds
Single covalent bond = 2 atoms sharing 1 pair of electrons H2
2 dots in an electron dot diagram represents this bond
A dash in a structural formula represents this bond
A molecular formula does NOT show this bond only the # of atoms
Halogens form these bonds in their diatomic molecules
Unshared pair = pair of valence electrons not shared in an electron dot diagram
*CP 8.1, PP 7-8 pg. 220

24. Double and Triple Covalent Bonds
Double covalent bond = involves 2 shared pairs of electrons
oxygen
Triple covalent bond = involves 3 shared pairs of electrons
nitrogen

26. Coordinate covalent bonds
The shared electron pair comes from one of the bonding atoms
CO – look at pg. 223
Polyatomic ion = a tightly bound group of atoms that has a positive or negative charge and behaves as a unit
NH4+
*c.p. 8.2, p.p pg. 225
A sign outside the chemistry hotel reads "Great Day Rates, Even Better NO3-'s"

28. Bond Dissociation Energies
Bond dissociation energy = the energy required to break the bond b/w 2 covalently bonded atoms
a large bond dissociation energy corresponds to a strong covalent bond
H2 = 435 kJ/mol, C-C single bond = 347 kJ/mol, C=C double bonds = 657 kJ/mol, and triple bonds = 908 kJ/mol

29. Resonance Structure
Resonance structure = a structure that occurs when it is possible to draw 2 or more valid electron dot structures that have the same # of electron pairs for a molecule or ion
Double-headed arrows are used to connect
Double bonds are usually shorter than single bonds but they are the same lengths b/c it is an avg. of the 2 structures
resonance

30. Exceptions to the Octet Rule
Cannot be satisfied in molecules whose total # of valence electrons is an odd #.
NO2, ClO2, NO
Sometimes w/ an even # as well (Fewer or more)
BF3
PCl5
SF6

31. *8.2 sect. assessment pg. 229*
*Exceptions to the octet rule: A resonance hybrid teacher DEMO*

32. Molecular Orbitals
Molecular orbitals = orbitals that apply to the entire molecule
just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole
Bonding orbital = a molecular orbital that can be occupied by 2 electrons of a covalent bond

33. Sigma Bonds
Sigma bonds = formed when 2 atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting 2 atomic nuclei
Atomic orbitals overlap end to end
Two s orbitals can combine to form a molecular orbital H2
Two p orbitals F2
The attractions b/w electrons and nuclei of two atoms overpower the repulsions b/w the 2 nuclei or b/w the 2 sets of electrons = covalent bond (stable molecule)

35. Pi Bonds
Pi bond = a covalent bond in which the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms
Orbitals overlap side by side
Atomic orbitals in pi bonding overlap less than in sigma bonding – weaker than sigma bonds
A typical double bond consists of 1 sigma & 1 pi bond, triple bond is 1 sigma & 2 pi bonds
In special cases, they form w/o any sigma bonds

37. VSEPR Theory
VSEPR theory (valence-shell electron-pair repulsion theory) = the repulsion b/w electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible
Explains the actual 3-D shapes of molecules

38. Hybrid Orbitals
Hybridization = several atomic orbitals mix to form the same total number of equivalent hybrid orbitals
Single bonds – one 2s and three 2p orbitals mix to form four sp3 hybrid orbitals
Double bonds – one 2s and two 2p orbitals mix to form three sp2 hybrid orbitals
Triple bonds – one 2s and 1 2p orbitals mix to form two sp hybrid orbitals

39. Electronic Geometry 2 electron densities – Linear
3 e- densities – Trigonal Planar
4 e- densities – Tetrahedral
5 e- densities – Trigonal Bipyramidal
6 e- densities – octahedral

40. Linear Atoms connected in a straight line
All molecules w/ 2 atoms and some w/ 3
180⁰ bond angle
HCl, CO2
Hybridization – sp
2 bonds/0 lone pairs

41. Trigonal Planar Triangular flat 120⁰ bond angle BCl3
Hybridization – sp2
3 bonds/0 lone pairs
Bent – 2 bonds/1 lone pair
118⁰
SO2

42. Tetrahedral 4 surfaces 109.5⁰ bond angle CH4 Hybridization – sp3
4 bonds/0 lone pairs
Pyramidal (trigonal pyramidal) – 3 bonds/1 lone pair
107⁰
NH3
Bent – 2 bonds/2 lone pairs
105⁰
H2O

43. Trigonal Bipyramidal 90, 120, and 180⁰ bond angles PF5
Hybridization – sp4
5 bonds/0 lone pairs
See-saw – 4 bonds/0 lone pairs
90, 120, and 180⁰
SF4
Tee-shaped – 3 bonds/2 pairs
90 and 180⁰
ClF3
Linear – 2 bonds/3 lone pairs
180⁰
XeF2

44. Octahedral 90 and 180⁰ bond angles SF6 Hybridization – sp5
6 bonds/0 lone pairs
Square pyramidal – 5 bonds/1 lone pair
90 and 180⁰
BrF5
Square Planar – 4 bonds/2 lone pairs
XeF4
8.3 sect. assessment pg

46. The VSEPR Model .. .. .. The Shapes of Some Simple ABn Molecules O S O
Linear
Bent
Trigonal
planar
Trigonal
pyramidal
AB6 F P F S F Cl F Xe
Students often confuse electron-domain geometry with molecular geometry.
You must stress that the molecular geometry is a consequence of the electron domain geometry.
The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.
T-shaped
Square
planar
Trigonal
bipyramidal
Octahedral

47. Bond Polarity Nonpolar covalent bonds = equal sharing of electrons
H2, O2, N2, Cl2
Polar covalent bond = unequal sharing of electrons
HCl, H2O
The more electronegative atom attracts electrons more strongly and gains a slightly negative charge, the less electronegative atom has a slightly positive charge

48. Table 8.3 pg. 238
C.P. 8.3, P.P pg. 239

49. Polar Molecules
Polar molecule = one end of the molecule is slightly negative, the other is slightly positive
Dipole = a molecule that has 2 poles w/opposite charges
The shape of a molecule and the polarity of its bonds together determine whether the molecule is polar or nonpolar
Equal and opposite directions arrows cancel = nonpolar
Arrows same direction = polar

51. Attractions b/w Molecules
Weaker than ionic or covalent bonds
Van der Waals forces = weak intermolecular attractions
1. Dipole interactions = intermolecular forces resulting from the attraction of oppositely charged regions of polar molecules
2. Dispersion forces = (weakest of all) attractions b/w molecules caused by the electron motion of one molecule affecting the electron motion on the other through electrical forces
Hydrogen bonds = 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 combo of this strongly polar bond and the lack of shielding effect in a hydrogen atom is responsible for the relative strength of hydrogen bonds (strongest of all intermolecluar attractions)

52. Table 8.4 pg. 244
*8.4 section assessment pg. 244