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?

Answers

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

Ch. 7 Ionic and Metallic Bonding

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*

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

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!"

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.

193 3-11

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.

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 mallTeen 1: Did you hear oxygen and magnesium got together?? Teen 2: OMg!

Your mama's so uglyYour mama's so ugly...even Fluorine won't bind to her

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 12-13 pg. 196, – 7.2 sect. assessment pg. 199 14-22

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-

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*

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

Ch. 8 Covalent Bonding

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*

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. 216 1-6

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

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

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. 9-12 pg. 225A sign outside the chemistry hotel reads "Great Day Rates, Even Better NO3

-'s"

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

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

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

*8.2 sect. assessment 13-22 pg. 229**Exceptions to the octet rule: A resonance hybrid teacher DEMO*

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

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)

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

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– http://

gold.chem.wwu.edu/sdgchem121/Docs/WWUDocs/geometry.pdf

– http://www.youtube.com/watch?v=i3FCHVlSZc4

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

Electronic Geometry

• 2 electron densities – Linear• 3 e- densities – Trigonal Planar• 4 e- densities – Tetrahedral• 5 e- densities – Trigonal Bipyramidal• 6 e- densities – octahedral

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

Trigonal Planar

• Triangular flat• 120⁰ bond angle• BCl3

• Hybridization – sp2

• 3 bonds/0 lone pairs– Bent – 2 bonds/1 lone pair• 118⁰• SO2

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

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

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• 90 and 180⁰• XeF4

• 8.3 sect. assessment pg. 236 23-29

......

The VSEPR Model

O OC

Linear

The Shapes of Some Simple ABn Molecules

O OS

BentO O

S

O

Trigonalplanar

FF

F

N

Trigonalpyramidal

T-shaped Squareplanar

F FCl

F

F F

Xe

F FF

F

FP

F

FTrigonal

bipyramidalOctahedral

FF

F

S

F

F

F

AB6

SO2

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

Table 8.3 pg. 238

• C.P. 8.3, P.P. 30-31 pg. 239

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

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)

• Table 8.4 pg. 244• *8.4 section assessment 32-38 pg. 244