Unit 2: Chemical BondingTopic-1: Verbal review in class1. Valence electrons (review)2. Lewis dot structures (review)3. Electronegativity (review)4. Cations[+] and Anions[-] (review)5. Octet Rule6. Bond energy graph (New)

Topic-2: Intra molecular chemical Bonds1. Metallic bonds 2. Covalent bonds3. Ionic Bonds4. SUMMARY: electronegativity difference and bond type prediction

AP Chemistry _ NotesDr. Chirie Sumanasekera

9/9/ 2019

Topic 1:

Review

Most electronegative (electron-greedy) atoms are in redLeast electronegative atoms are in yellow

Scale of Valence electron sharing

Non-polarCovalent bonds

PolarCovalent bonds

Ionic bonds

Equal sharing Unequal sharing

Large difference in electronegativity

Similar electronegativity

Increasing bond energy

Covalent and ionic bonds are formed by a continuum of electronegativity difference/bond energy between the atoms forming the bond

Bond length and Bond P. Energy graphhttps://www.bing.com/videos/search?q=potential+energy+and+inter+nuclear+distance+graph&&view=detail&mid=5FE777F2D46C88E65F535FE777F2D46C88E65F53&rvsmid=86F3CD8E274BB882AF9F86F3CD8E274BB882AF9F&FORM=VDQVAP

The bond is formed at this energy level, where the repulsive forces and the attractive forces are equal

1. too far for attractive forces2. only attraction3. Attraction is equal to repulsion4. Only repulsion

INTRA-Molecular BONDS(bonds inside molecules and ionic compounds or

bonds formed between atoms)

1. Metallic bonds

• How metallic bonds are formed • Properties of metallic compounds • Metal alloys and their structure: Substitution vs.

intersitial alloys

1. Metallic Bonds

• The attraction between metal cations and the surrounding sea of delocalized valence electrons is called metallic bond.

Delocalized valence electron (mobile)Positively charged ions- held in a rigid lattice

Metal atomic structure:

Inside metals, positively charged ions (cations) swim in a sea of delocalized valence electrons. These electrons are shared by all neighboring atoms and are free to move around.

Properties of metallic bonds• Metals form compact, solids at room temperature

(Except Mercury /Hg)

• Metals are good electrical conductors as they have

readily mobile electrons to produce a current

• The shared-electron mobility is also why metals are

malleable (hammered into shapes) and ductile

(stretched into wires).

• Color: Shiny- metallic

• Dense

• High melting point and boiling point

• Conduct heat

Alloys of metals

Examples of alloys: 1. Bronze: Cu + Tin + Al/Mn/Zn

sterling silver,2. Cast iron: Fe + C + Si3. Steel: Fe + C

Metal alloys can be of two types:1. Substitution alloy = another metal atom of

equal size is substituted in the space of a former metal atom (made up of 2 or more metals)

1. Interstitial alloy: made up of metal atoms and another smaller metal or non metal atom that is in the space between metal atoms.

Q: Why is steel stronger than cast iron? Explain your answer

What is a metal Alloy?A mixture made of at least one metal plus a metal or non metal.

2. Covalent bonds• How Ionic covalent bonds are formed • Properties of covalent compounds • Types of covalent bonds (Polar, non polar, double,

single and triple)• Lewis dot structures and predicting structure of

compounds (what is a coordinate covalent bond?)• Predicting the polarity of covalent compounds based

on electronegativity of various atoms

2. Covalent bonds

• What is a covalent bond? Covalent bonds are intra-molecular bonds formed when atoms SHARE electrons. All compounds that make molecules are formed by covalent bonding of the atoms within.

• In covalent bonds, electrons are usually shared between 2 or more non-metal atoms so that each can obtain the noble gas electron configuration.

• There can be single, double or triple covalent bonds between two atoms

• Based on the electronegativity values of each atom in a covalent bond, the resulting molecule may have a dipole (polarity) based on negative and positive charged regions of the same molecule. [Polarity is explained with the symbol delta = d - or d +]

• Covalent bonds and ionic bonds exist in a spectrum (continuum) going from zero to extreme electronegativity

Properties of covalent-Bonds

• Bond strength varies depending on type of bond and atoms involved in bonding

• Can be found in solid, liquid or gas state depending on the compound

• Poor electrical conductors (as they do not have readily mobile electrons to

produce a current)

• Can be colorless or have colors

• Low melting point and boiling point

• Do not conduct heat well

Types of Covalent Bonds1. Single bonds [ formed by sharing 1 electron-pair = 2 electrons between 2 atoms]2. Double bonds [ formed by sharing 2 electron-pairs = 4 electrons between 2 atoms]3. Triple bonds [ formed by sharing 3 electron-pair = 6 electrons between 2 atoms]

Note: • Each bond is represented as a straight line. A single bond will have one line, double

will have 2 lines and a triple bond will have 3 lines.• Triple bonds are the strongest and single bonds are the weakest and easiest to break• Bond distance is shortest in triple bonds and longest in single bonds. See figure below:

Draw the Lewis dot structures of the following covalent compounds:

1. HCl

2. CO2 (carbon dioxide)

3. N2 (nitrogen gas)

4. H2 (hydrogen gas)

5. O2 (oxygen gas)

Examples of Covalent Bonds

NN2 N N N NN+

HCl H Cl+ H Cl H ClSigma (Single) bond

Double bond

Triple bond

O C O C OOCO2 2+C O

Covalent Bonds of the Water molecule

O

O HH

H2O +2 H

Non bonding lone electron pair

HOH

Water: Polar Covalent Bonds

OHH

OHH

d +d +

d - Electron Density shifts towards the more electronegative oxygen atomO

HH• Oxygen is much more electronegative than Hydrogen, so Oxygen pulls the shared

electrons towards itself

• This greater electronegativity of oxygen compared to hydrogen in the covalent bonds of water molecules result in a dipole or polarity. The result is a shifting of the electron density towards Oxygen.

• H gets a small positive charge Delta (d +) and Oxygen gets a negative charge (d -)as electrons are less equally shared between O and H.

Non-polar covalent bonds

NN2 N N N NN+

• If a molecule if formed by the same kind of atom or atoms with similar electronegativity then the covalent bond is non-polar (Ex: O2, N2, Cl2, O3, Br2)

• Carbon and H makes non polar covalent molecules such as hydrocarbons(CH4 , C2H6 , C3H8 and so on.)

• DRAW C2H6 (not C2H8)

Coordinate covalent bonds

In typical covalent bond, each atom in the bond provides an electron for a single bond. But in coordinate covalent bonds this is not the case. Here, one atom contributes both of the shared electrons for one covalent bond.

OOC OCEx: CO

Carbon monoxide

C

OC

Both these electrons come from Oxygen

ReviewMatch the correct numbers with the questions on the right

H

1. Triple bond:___________

2. Double bond:___________

3. Single bond:___________

4. Unpaired electrons:__________

5. Region/s with highest electron density:__________

6. Region/s with lowest electron density:__________

7. Region/s that is d +

:__________

8. Region/s that is d - :__________

C

H

C

Cl

F

H1

CCC

O

C

2 3 4 5

6

7

8

9

12

10

11

H

1. Covalent Molecules: formed by covalent bonding (electron sharing) between atoms

2. Ionic (bonded) compounds: Do not form molecules. They are formed by atoms that stick close together due to opposite charge (electrostatic) – attractions caused by electron exchange.

Review Question: How are covalent compounds (molecules) different from Ionic compounds?

3. Ionic Bonds:

• How Ionic bonds are formed• Properties of ionic compounds• Lattice energy and structure of ionic compounds• Predicting the formulas of an ionic compounds

• Ions with opposite charges are mutually attracted to each other. • Attractions between oppositely charged ions are called electrostatic

attractions. • Ionic bonds are electrostatic attractions that occur between ions with

opposite charge.• Ionic bonds always occur between non-metal ions (Anions /-) and metal

ions (Cations/+).• All ionic compounds are electrically neutral because their total positive

and total negative charges add up to zero.• Ionic compounds arrange in a repeating pattern called a crystal lattice.

(like a lattice pie!)• The energy needed to break this lattice is called Lattice energy.• When the lattice is broken, the ions will become atoms with zero charge

and the ionic bonds will be broken.

3. Ionic bonds

Properties of Ionic bonds

• Form crystalline solids at room temperature• Generally have high melting points• When dissolved in water, can conduct electricity and

form electrolytes.

Ionic bond Ex: Table salt, NaCl

Lewis dot structure for NaCl:

• Na is in group 1A = donate the valence electron 1 e -• Cl is in Group 7A = accept this valence 1 e -• Na donates 1 electron and become an Na+ ion and Cl will accept 1 electron

and become a Cl- ion

Ex: Write the net ionic equation for the formation of Na+ and Cl– ions:

Na – 1e– Na+ + 1e– (1 electron donated by sodium)

Cl + 1e– Cl– (Sodium electron accepted by Cl)

Predicting the formulas of ionic compoundsCrisscross-method of predicting Ionic compound formulas:

Step-1: Write the name of the atoms side by side:

Step-2: Find out the number of valence electrons for each atom and the Charges made when they ionize (become ions). Ex: +2 or -1 etc. Write the Charge above the atoms (super-scrips):

Step-3: Now swap the charges as shown by the arrows:

Step-4: Then write the swapped ionic-charges as coefficient in front of the Ion:

Note Ones are not shown when writing formulas:

This means, one Cl-ion combines with 3 Al-ions in the ionic compound. This will become clear to you once your draw the Lewis dot structures of the atoms and the compound.

Al Cl

1Al+3 3Cl-1

Al+3 Cl-1Al+3 Cl-1

Al+3 + 3Cl-1 Final answer

Predicting the formulas of ionic compounds

This means one Cl-ion combines with 3 Al-ions in the ionic compound. This will become clear to you once your draw the Lewis dot structures of the atoms and the compound.

Step-5: Now draw the combined electron dot structure for the compound and determine how many of each ion is found in a unit of the compound.

Since each Cl atom needs only 1 electron and Al atoms has 3 valence electrons to donate, 3 Cl ions will come bine with a single Al atom to take its 3 valence electrons and complete their octets! - -

Al Cl

Al ClCl

Cl

Cl

AlAl+3

Cl-

Cl-

Cl-

OR

-

+33

Predict the formulas of the ionic compounds below using the criss-cross method:

5. Mg and Cl

6. Mg and S

7. Mg and P

8. Mg and C

1. Na and Cl

2. Na and S

3. Na and P

4. Na and C

9. Al and Cl

10. Al and S

11. Al and P

12. Al and C

13. C and Cl

14. C and S

15. C and P

16. Mg and C

4. Electronegativity difference and bond type

Electronegativity difference & bond types SUMMARY:

Polar covalent

bonds0.4 - 2.0

Ionic bonds

2.0 and greaterElectronegativity difference:

Non-polarcovalent

bonds0.0-0.4

Increasing electronegativity

Examples: H H H Cl

H F

Na+ Cl-

Unequal electron sharing

Equal electron sharing

**Bond length decreases when number of bonds increase**Bond Strength increases when number of bonds increase

Properties of Chemical Bonds

Property Meaning1 Bond Length Distance between the nuclei of two bonded atoms

( unit = picometers =pm)2 Bond angle Angles of any two bonds around an atom3 Bond energy

(bond dissociation energy)Energy required to break a bond.( unit = Kilo Joules per mole = KJ/mol)

Calculate electronegativity differences:

1. F:_______

2. H:______

3. C:______

4. H-F:_____

5. C-C:____

6. C-H:_____