Chemical Bonds

Mutual attraction that binds atoms together to form compounds

Types of Chemical Bonds

• Ionic bond = bond that results from attraction between oppositely charged ions (transfer of electrons)

• Covalent bond = bond resulting from the sharing of electrons

Metallic Bonds

= metal atoms in a delocalized cloud of electrons• Remember: metals like to give up electrons,

so no atom “wants” the free e-

Bonding happens on a spectrum

• Bonds are rarely purely ionic or purely covalent

• Remember Electronegativity (the ability of atoms to attract electrons)– Comparing the electronegativities of atoms

involved in bond, can determine whether bond is ionic or covalent

see page 161 PT for electronegativity values

Electronegativity difference above 1.7 = ionic

• Ex. Cs and F

• Cs = 0.7• F = 3.3

Electronegativity difference of 1.7 or less = covalent

• Polar-covalent (0.3-1.7), – ex. H2O: H-2.1, O-3.5, diff = 1.4

• Non-polar covalent (under 0.3)– Ex. Bonds between atoms of the same element

are always non-polar (purely) covalent, ex. H2, O2, N2

Covalent Bonding

Forms molecular compounds that consist of molecules

Molecular Compounds

• Molecule = a neutral group of atoms that are held together by covalent bonds

• Chemical formula (molecular formula) = indicates the number of atoms of each element in a compound– H2O

– O2

– CO2

– HCl– C6H12O6

Diatomic molecule = made up of 2 atoms of the same element

Diatomic Molecules (Hydrogen and the magic 7)

• Some elements always exist as diatomic molecules

Empirical Formula

• The “reduced” form of a compound• Has the lowest ratio

• Molecular formula = C6H12O6

• Empirical formula = CH2O

Potential Energy and Bonding

• Nature favors covalent bonding– Atoms have lower PE when they are bonded– As atoms near each other, their charged particles

interact• Nucleus is attracted to e- + - PE↓ • e- repel each other - - PE↑• 2 nuclei repel each other + + PE↑

Characteristics of Covalent Bonds

• Bond length = avg. distance btw 2 bonded atoms

• Bond energy = E required to break the bond

Covalent Bond = overlap of orbitals where shared electrons live

How are they related?

Octet Rule

• Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet (8) electrons in its highest occupied energy level (valence)

• F: ↑↓ ↑↓ ↑↓ ↑↓ ↑__ 1s 2s 2p

• F: ↑↓ ↑↓ ↑↓ ↑↓ __↓

Bonding electron pair in overlapping orbitals

Example

1sH: ↑__ 1s 2s 2p 3s 3pCl: ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ __↓

Bonding electron pair in overlapping orbitals

Exceptions to the Octet Rule

• Hydrogen (1 valence) and Helium (2 valence) experience stability with only 2 electrons in the 1s-orbital (it is their highest and ONLY energy level)

• Boron (3 valence) tends to form stable compounds with 6 valence e-

– Ex. BF3

• Other elements can be surrounded by more than 8 e- when they bond with highly electronegative elements (like halogens) – This bonding will involve d orbitals In addition to s and p– Ex. PF5 and SF6

Electron Dot Notation

1 X Na2 X Mg3 X B4 X C5 X N6 X O7 X F8 X Ne

# valence e- electron-dot notation example

Lewis Structures = formulas in which atomic symbols represent nuclei and inner shell electrons; only valence electrons are drawn

• Use electron-dot notation to represent molecules– Examples:

Structural Formulas indicate the kind, number, arrangement, and bonds, but not the unshared pairs of electrons

Types of Covalent Bonds

• Single bonds = result from the sharing of 1 electron pair

• Double bonds = result from the sharing of 2 electron pairs

• Triple bonds = result from the sharing of 3 electron pairs

Bond Length and Bond Energy

• REMEMBER: bond energy is energy required to break the bond

• Shorter bonds have higher energies and are harder to break

single bonds double bonds triple bonds

Shorter length

Higher energy

Molecular Geometry

= the 3-dimensional arrangement of a molecule’s atoms in space

Geometry and bond polarity will determine molecular polarity and intermolecular forces.

Forces of attraction between molecules

Intramolecular Forces

• Forces between atoms or ions that keep compounds together

– Ionic attraction– Sharing electrons

Intermolecular Forces

1. Dipole-Dipole forces = attractive forces between polar moleculesHydrogen-bonds (H-bonds) = type of D-D force

btw H and certain other elements H-N, H-O, or H-F

2. Ion-Dipole forces = attraction btw an ion and a polar molecule

Intermolecular Forces (cont)

3. London Dispersion forces = attractive forces resulting from temporary dipoles induced by ions

Intermolecular Forces (cont)

Naming Covalent Compounds

• Use the prefixes based on the number of each type of atom and suffix –ide for 2nd element

1 = mono- CO = carbon monoxide2 = di- CO2 = carbon dioxide3 =tri- BF3 = boron trifluoride4 = tetra- CCl4 = carbon tetrachloride5 = penta- N2O5 = dinitrogen pentoxide6 = hexa- B3F6 = triboron hexafluoride7 = hepta-8 = octa-9 = nona-10 = deca

Only used in naming organic compounds

Naming Covalent Compounds

• Exceptions: covalent compounds containing hydrogen, e.g. – H2S (hydrogen sulfide)

– CH4 (methane)

– NH3 (ammonia)

Organic Compounds

• Compounds containing carbon– Hydrocarbons = carbon chain + hydrogen

– Carbohydrates = carbon + hydrogen + oxygen

Propane

Glucose

Naming Hydrocarbons

Consist of chains of carbon atoms

Name = prefix + -ane

# Carbon atoms prefix

1 Meth-

2 Eth-

3 Prop-

4 But-

5 Pent-

6 Hex-

7 Hept-

8 Oct-

9 Non-

10 Dec-

Drawing Hydrocarbons

Ionic Compounds

= Compounds consisting of a positive ion and a negative ion held together by the attraction of opposite electrical charges

The overall charge of an ionic compound = 0The + and - ions cancel each other

Formula Units

• Molecules are the basic unit of covalent compounds (a molecule of water, H2O)

• The smallest unit of an ionic compound is the formula unit (a formula unit of sodium chloride, NaCl)

Crystal Lattice

= an orderly arrangement of ionsAttractive forces

btw opposite ionsbtw nuclei and e- of adjacent ions

Repulsive forces btw like ionsbtw e- and e- btw nucleus and nucleus

Lattice Energy and Bond Strength

To compare bond strength in ionic compounds, we compare lattice energy

Related to the energy in the bond = how tightly the ions in a crystal are held togetherRemember: higher energy bonds are harder to break

Molecular (covalent) vs. Ionic

• Covalent and ionic bonds are strong attractions that hold the atoms in a compound together

Molecules in Covalent CompoundsHeld together by intermolecular forces

Formula Units in Ionic Compounds Held together by attractive forces in a crystalline latticeSTRONGER

Molecular v. Ionic Properties

Molecular• Lower melting points(liquids and gases)

Ionic• Higher melting points(solids)• Hard and brittle• Conduct electricity as

liquids or when dissolved in water

• Many are soluble in water

• Some aren’t. WHY?

Types of Ions

• Monatomic = ions that consist of one element

Cl- = chlorideBr- = bromideNa+

Mg2+

Al3+

• Polyatomic = ions that consist of more than one type of atom

(CO3)2- = carbonate

(PO4)3- = phosphate

(C2H3O2)1- = acetate

(ClO3)1- chlorate

(NO3)1- = nitrate

Naming Binary (2 elements) Ionic Compounds

Cation (metal) comes 1st

– Element nameAnion (nonmetal) comes 2nd

– Element root + -ide

Naming Ionic Compounds that Contain Polyatomic Ions

• Cation (metal)– Element name

• Polyatomic ion– Ion name

Stock System for Transition Metals

• Some transition metals (Group 3-12) can form more than one type of cation – Ex. Copper can form +1, +2, or +3

• When writing the name, use a Roman numeral to indicate the charge of the transition metal ion– Ex. Copper (I) chloride, copper (II) chloride, etc.

Writing Formulas for Ionic Compounds

1. Symbol for cation + symbol/formula for cation

2. Write the charges as super scripts3. Cross the numbers down to subscripts

Naming Acids

• Binary Acids = consist of 2 elementsHydrogen + a halogen• HF = hydrofluoric acid• HCl = hydrochloric acid• HBr = hydrobromic acid• HI = hydroiodic acid

Naming Acids

• Oxyacids = consist of H, O, and another nonmetal

H2SO4 = sulfuric acidH2SO3 = sulfurous acid

H2CO3 = carbonic acid

HNO3 = nitric acidHNO2 = nitrous acid

HC2H3O2 = acetic acid

HClO4 = perchlorous acidHClO3 = chloric acidHClO2 = chlorous acidHClO = hypochlorous acid

Assigning Oxidation Numbers

• The oxidation # for any uncombined element or diatomic molecule is ZERO

• The oxidation # for any monatomic ion is its CHARGE

• H is usually +1, unless it’s combined with a metal to form a metal hydride then it’s -1

• O is usually -2, unless it’s a peroxide, then it’s -1

Ex. Ba+2, K+1

Ex. Zn, Cu, O2, N2, Cl2

HCl (+1), NaH (-1)

MgO (-2), H2O2 (-1)

• In binary covalent compounds (nonmetal + nonmetal) the positive one is first and the negative one is second

• The sum of the oxidation numbers for all atoms in a neutral compound is ZERO

• The sum of the oxidation numbers in a polyatomic ion is equal to the CHARGE of the polyatomic ion

Assigning Oxidation Numbers

Practice

• H2

• CaCl2

• KClO4