Chemical Bonding

and MolecularStructure

Cartoon courtesy of NearingZero.net

BONDS BETWEEN ATOMS (Intramolecular) Attraction of Valence electrons between nuclei

based on their electronegativityWhen atoms combine, there is a tug of war over

the valence electrons Electrons are either gained, lost or shared in order

to achieve a complete OCTET OBTAINING A NOBLE GAS CONFIGURATION (STABLE)

Bonded atoms are at a lower energy condition (lower PE)

Ionic bonds – transfer of electrons Covalent bonds – sharing of electrons

Bond Stability• WHEN A BOND IS FORMED energy is RELEASED (exothermic)

and the resulting compound has LESS energy and is more STABLE. – A + B → AB + 400 Joules (Energy on the product side)

• WHEN A BOND IS BROKEN energy is ABSORBED (endothermic).The resulting substances are in a high energy state and are UNSTABLE– AB + 400 Joules → A + B (Energy on the reactant side)

• Low energy states are more stable than high energy states therefore the greater the energy released the more stable the bond (see table I)

Pneumonic: B.A.R.F. Break Bonds-Absorb EnergyRelease Energy- Form Bonds)

Types of Bonds Metal/nonmetal ionic Nonmetal/nonmetal covalent Metals only metallic

• ELECTRONEGATIVITY difference (END) can be used to predict the type of bonding & character of the substance

Nonpolar polar ionic

covalent covalent

• The greater the END the more polar or ionic the bond

Ionic Bonds (salts)• Metal plus Non-

metal • composed of

metals bonded to nonmetals to form ionic crystals (lattice).

Examples:

Ba2+S2- Mg2+Cl-2Na+2O2-K+I-

Representation of Components in an Ionic Solid

Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

• Metals LOSE electrons forming cations, nonmetals lose electrons forming anion. Two ions are produced with OPPOSITE charges, therefore they attract. This is referred to as an ELECTROSTATIC ATTRACTIONwhich produces an electrically neutral substance.

• END is greater than 1.7 (use Table S)• The Octet Rule Diagram– (Bohr Diagram)

Na 2-8-1 Cl 2-8-7 Na+ 2-8 Cl- 2-8-8This transfer forms ions, each with an octet

The Formation of Sodium Chloride

Cl-Na+

• The resulting ions come together due to electrostatic attraction (opposites attract):

• The net charge on the compound must equal zero• Electronegativity of 1.7 or greater indicates an

ionic bond.

• NaCl is the chemical formula for sodium chloride.

l

PROPERTIES OF IONIC COMPOUNDShttps://www.youtube.com/watch?v=TxHi5FtMYKk

HIGH melting and boiling points Hard, crystalline, solid substances (at room temperature)

ie. NaCl (s)

Do not conduct electricity UNLESS dissolved in water (ex. NaCl (aq) )

Molten/ liquid/melted forms can conduct electricity (ex NaCl(l)) due to the MOBILE charged particles called IONS

Summary of Propertiesof Ionic Compounds

Structure: Crystalline solidsMelting point: Generally highBoiling Point: Generally highElectrical Conductivity:

Excellent conductors, molten and aqueous

Non-conductors in solid stateSolubility in water:

Generally soluble

Electron Dot Notation

Represents the element surrounded

by its valence electrons

LEWIS DOT DIAGRAMS OF IONIC COMPOUNDS

• Illustrate the transfer of electrons between metals and non-metals following the octet rule.

• Metals lose ELECTRONS WHEREAS Non-metals GAIN electrons

• Use [brackets] around the ion and put charge in the upper right

(Metals) (nonmetals)

• Ionic – show transfer of e-

• Your turn - Ionic Bond Practice

• The ion charges CANCEL out (add up to zero) in compounds

Ex. MgCl2

Metallic Bonding Metallic Bonding and the Properties of Metal video https://www.youtube.com/watch?v=Oagr9xMAmfY

1. Giant crystalline structure of metal only; (+) ions are in FIXED patterns and valence electrons move about freely –“sea of mobile electrons”

• Made up of metals ONLY

2. Look for a Single Metal - ex. Cu, Mg, Fe

Metallic Bonds and Metallic Properties• A force can change the shape of a metal. A force can shatter an ionic crystal.

• A metal rod can be forced through a narrow opening in a die to produce wire. a) As this occurs, the metal changes shape but remains in one piece. b) If an ionic crystal were forced through the die, it would shatter.

What causes the ionic crystal to break apart? The repulsion between the like charges of the ions

3. Properties of Metals• Good conductors of heat and

electricity• Solids• Malleable and Ductile• High melting and boiling points

• Metal atoms are arranged in very compact and orderly patterns.Crystalline Structure of Metals

ALLOY – molten mixture of two or more elements, one being a metal• Ex. Steel (Fe, Cr, C, Ni)

and Sterling silver (Ag and Cu)

• Alloys are important because their properties are often superior to those of their component elements.

• The most important alloys today are steels. Steels have a wide range of useful properties, such as corrosion resistance, ductility, hardness, and toughness.

Covalent Bonds: Molecular Substances1. A bond between two or more NONMETALS• This happens when non-metals SHARE their valence

electrons to achieve a full octet.• May exist as solids, liquids or GASES (depending on

intermolecular attractions)

A compound composed of molecules is called a molecular compound. Water and carbon monoxide are molecular compounds. The bonds between H and O are covalent.

2. PROPERTIES of Molecular (COVALENT) substances

– LOW melting & boiling points– Poor conductors – Soft solids ex. H2O, CO2, NH3, glucose

COVALENT NETWORK solids

•LARGE molecules (MACROMOLECULES)•Hard, poor conductors, HIGH mp/bp.

•ex. DIAMOND, GRAPHITE, SiC, SiO2 (quartz) • used for abrasive & cutting tools

3. Bonds can be SINGLE, DOUBLE or TRIPLE

• If one electron is needed to achieve an octet – SINGLE ex. HCl

• If two electrons are needed -DOUBLE ex. CO2

• If three electrons are needed -TRIPLE ex. N2

Shows how valence electrons are arranged among atoms in a molecule.Reflects central idea that stability of a compound relates to noble gas electron configuration.

Lewis Structures

Drawing the Lewis Dot structure of a molecular compound:

• Remember that these are not ionic compounds and do not need brackets or charges because the atoms SHARE electrons.

***Note that the oxygen needs 2 electrons and each of the hydrogen need 1 electron. So they will share 4 electrons creating 2 bonds.

• Each line represents 2 electrons that are shared

CH

H

H

Cl

..

....

Completing a Lewis Structure -CH3Cl

Add up available valence electrons:C = 4, H = (3)(1), Cl = 7 Total = 14

Join peripheral atoms to the central atom with electron pairs.

Complete octets on atoms other than hydrogen with remaining electrons

Make carbon the central atom

..

..

..

Single Covalent BondsElectron (Lewis) Dot Diagrams

The methane molecule has four unshared

pair of electrons.

The ammonia

molecule has

one unshared

pair of

electrons.

Covalent Bonds for the Diatomic ElementsElectron (Lewis) Dot Diagrams

Two atoms

held together

by sharing a

pair of

electrons are

joined.

Double and Triple Covalent Bonds

The oxygen

molecule is an

example of

a diatomic

molecule.

The carbon

dioxide molecule

is an example of

a triatomic

molecule.

Lewis Dot Structures• See guide pg 4 (read directions)

Types of Covalent BondsNON-POLAR COVALENT BONDS

• EQUAL SHARING OF VALENCE ELECTRONS• END between nuclei is ZERO -.4• ex. “HOFBrINCl” • Diagram Cl2

POLAR COVALENT BONDS• UNEQUAL sharing of valence electrons• END between nuclei is less than 1.7 but > 0.4• In a polar bond the atom with the higher

ELECTRONEGATIVITY is slightly negative whereas the atom with the lower EN is slightly positive– The greater the END, the more polar the bond and the

greater the IONIC CHARACTER – The lower the END, the less polar and WEAKER the ionic

character of the bond• Diagram: HCl

Coordinate Covalent Bonds• A coordinate covalent bond is a covalent bond in which

one atom contributes both bonding electrons.– involves the POLYATOMIC ions (Table E)– one atom shares NO electrons whereas the other

contributes both• diagram: ex. NH4

+ ex. H30+

In a structural formula, you can show coordinate covalent bonds

as arrows that point from the atom donating the pair of

electrons to the atom receiving them.

Covalent –vs – Ionic Bonds

Bond Polarity

Predicting the polarity of the bond

0.0 - 0.3 Non-polar Covalent

0.4 – 1.0 Moderately polar covalent

1.0 – 1.6 Very polar covalent

1.7 or greater ionic

VSEPR Model - (Valence Shell Electron Pair Repulsion)

The structure around a given atom is determined principallyby minimizing electron pair repulsions.States that because electron pairs repel, molecules adjust their shapes so that valence electrons are placed as far apart as possible

• Angle between any two bonds

ex. A _ B…..No angle (Linear)

TYPES OF MOLECULES (POLARITY) DUE TO SHAPE – ONLY covalent bonds

POLAR MOLECULES• Also referred to as a dipole (2 poles)• Molecules in which there are polar covalent

bonds and have an ASYMMETRICAL charge distribution

• UNEQUAL distribution of electrons (Winner in Tug of War)– Ex. H2O, NH3, HCl

δ+ δ−

FHmolecule HF

When two unlike atoms are covalently bonded, the shared electrons are strongly attracted to the atom of greater electronegativity.

• Nuclei with the greater EN gets a (-) charge, whereas the nuclei with the lower EN receives the (+) charge - called a polar DIPOLE (NONSYMMETRICAL)– Shape: bent or pyramidal & has 0-1 number of

Lines of symmetry• Examples:• H2O NH3

bentpyramidal

NONPOLAR MOLECULES• A molecule with either nonpolar bonds or polar bonds

and is SYMMETRICAL• Electrons are EQUALLY distributed -This gives an

EQUAL distribution of the electron charge (No winner in tug of war)

• Shape: linear or tetrahedral shape: 2 or more number of Lines of symmetry

• ex. All “HOFBrINCl” (gases)• Ex. CH4 CO2

tetrahedral

linearContains 2 polar bonds & is linear therefore bond polarities cancel out

INTERMOLECULAR FORCES (IMF’s)

• Groups of atoms covalently bonded in a molecule which may be ATTRACTED to similar molecules or ions

• Attractive Forces BETWEEN molecules, NOT ATOMS

• Stronger the force, the higher the BP/MP

Hydrogen Bonds • Formed b/w molecules when hydrogen is bonded to an

element of small atomic radius and high EN. (FLUORINE, NITROGEN, & OXYGEN) – ex. H2O, HF & NH3

• STRONGEST force of all intermolecular attractions • Accounts for the high BOILING POINT OF WATER

– EX. H2O

Hydrogen Bonding in Water

• Hydrogen Bonding between Ammonia and Water

Hydrogen Bonding

Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests.

Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen

Water in the Liquid State

Surface tension makes it possible for some insects, such as this water strider, to walk on water. Water molecules at the surface of the water drop above cannot form hydrogen bonds with air molecules, so they are drawn into the body of the liquid, producing surface tension.

Dipole-Dipole Interactions

• Attractive Forces between polar molecules due to charges: opposites attract

Van der Waals Forces or London Dispersion forces• WEAKEST of all IMAF’s -occurs between all

non-polar molecules making it possible for small nonpolar molecules (H, N, O) exist in a liquid and solid phase under conditions of low temp and high pressure

• Forces increase with increasing mass and with decreasing distance /w the molecules (ex. Group 17)

• The effect of molecular size accounts for the increasing boiling points of a series of hydrocarbons

Forces increase as phase changes g,l,s

Molecule-Ion Attraction • Attraction of a polar molecule (water, alcohol, ammonia)

for the ions in a crystal lattice (salts) –opposite charges ATTRACTING forming a solution (aq) and DISSOLVINGthe ionic compound into solution

• Orienting of water molecules around ions is called “hydration of the ions”

• Diagram: NaCl(aq)

Molecule-Ion Attraction

Dissolution of sodium Chloride

Relative magnitudes of forcesThe types of bonding forces vary in their strength as measured by average bond energy.

Covalent bonds (400 kcal)

Hydrogen bonding (12-16 kcal )

Dipole-dipole interactions (2-0.5 kcal)

Van der waals forces (less than 1 kcal)

Strongest

Weakest

Intermolecular Attractions SummaryHydrogen Bonding

Van der waals forces

Molecule-ion attraction

Dipole interactions

Type ofCompound involved

Polar molecules with a H bonded to a highly EN atom

Non-polarMolecules or noble gases

Ionic solids (NaCl) in polar solventsH2O (Aq)

Polar molecules

Type of attraction

(+) H side of one molecule to the (-) end of the thehigh EN atom------ (H20)

Weak forcesForces asMolecular mass

as phase changes g,l,s

Opposites attract(+) pole attracted to anion (-) pole attracted cation

Attraction between oppositely charged regions of neighboring molecules.

Examples of compounds affected

H20 (accounts for high bp)NH3

FON w/ H

Group 17 G l sGroup 18

NaCl (aq) HClHBr

FORMING SOLUTIONS BETWEEN MOLECULES

• “likes dissolving likes”• depends on the forces of attraction• the POLAR bear fell in WATER (H2O) and caught AMMONIA (NH3)

– Soluble = DISSOLVES insoluble = doesn’t dissolve– Solute –gets dissolved solvent = does the dissolving

Solute type nonpolar solvent

(soap)

polar solvent (water)

Nonpolar (butter)

Polar (sugar)

Ionic (salt) insoluble

soluble

soluble

soluble

insoluble

insoluble

Practice1.Tell if the following pairs are soluble in each other: write soluble or insoluble

• NaCl and H2O _____________ CCl4 and I2 ____________

• LiF and CCl4 _____________ HCl and CH4 ______________

• CH4 and CCl4 _____________

2. Explain why gasoline (solute) on the hands cannot be removed by water

(solvent) alone? Use terms such as soluble, insoluble, likes dissolving likes, etc

to justify your answer.

3. Hairspray can remove ink stains from clothing but water alone cannot. Explain why?

(I & P) soluble

(NP & NP) soluble

(P & NP) insoluble(I & NP) insoluble

Water is polar and gasoline is nonpolar. They are insoluble and will not dissolve each other because likes dissolve likes.

Water is polar. Hairspray and ink are both nonpolar (soluble with each other). Likes dissolve likes.

(NP & NP) soluble