BONDING General Rule of Thumb: metal + nonmetal = ionic polyatomic ion + metal or polyatomic ion = ionic (both) nonmetal + nonmetal(s) = covalent.

BONDING

General Rule of Thumb:

metal + nonmetal = ionic

polyatomic ion + metal or polyatomic ion = ionic (both)

nonmetal + nonmetal(s) = covalent

Ionic BondsIsn’t it ionic that opposites attract?

Valence Electrons

Knowing electron configurations is important because the number of valence electrons determines the chemical properties of an element.

Valence Electrons: The e- in the highest occupied energy level of an element’s atoms.

Valence Electrons

All elements in a particular group or family have the same number of valence electrons (and this number is equal to the group number of that element)

Examples: Group 1 elements (Na, K, Li, H): 1 valence electron. Group 2 elements (Mg, Ca, Be): 2 valence electrons. Group 17 elements (Cl, F, Br): 7 valence electrons.

Lewis Structures

Electron dot structures show the valence electrons as dots around the element’s symbol:

Li B Si N O F Ne

Lewis Structures

Electron dot structures show the valence electrons as dots around the element’s symbol:

Li B Si N O F Ne

Octet Rule

Noble gas atoms are very stable; they have stable electron configurations. In forming compounds, atoms make adjustments to achieve the lowest possible (or most stable) energy.

Octet rule: atoms react by changing the number of electrons so as to acquire the stable electron structure of a noble gas.

Octet Rule

Atoms of METALS obey this rule by losing electrons.

Na: Na+: Atoms of NONMETALS obey this rule by

gaining electrons. Cl: Cl-: Transition metals are exceptions to this

rule. Example: silver (Ag) By losing one electron, it acquires a

relatively stable configuration with its 4d sublevel filled (pseudo noble-gas)

Octet Rule


Na: 1s22s22p63s1

Na+:1s22s22p6 = same # electrons as Ne

Atoms of NONMETALS obey this rule by gaining electrons.

Cl: Cl-: Transition metals are exceptions to this rule. Example: silver (Ag) By losing one electron, it acquires a relatively

stable configuration with its 4d sublevel filled (pseudo noble-gas)

Octet Rule


Na: 1s22s22p63s1



Cl: 1s22s22p63s23p5

Cl-: 1s22s22p63s23p6

Transition metals are exceptions to this rule. Example: silver (Ag) By losing one electron, it acquires a relatively


Octet Rule


Na: 1s22s22p63s1



Cl: 1s22s22p63s23p5

Cl-: 1s22s22p63s23p6



Octet Rule


Na: 1s22s22p63s1



Cl: 1s22s22p63s23p5

Cl-: 1s22s22p63s23p6



Ionic Bonds

Anions and cations have opposite charges; they attract one another by electrostatic forces (IONIC BONDS)

Ionic Bonds Ionic compounds are electrically

neutral groups of ions joined together by electrostatic forces. (also known as salts) the positive charges of the cations

must equal the negative charges of the anions.

use electron dot structures to predict the ratios in which different cations and anions will combine.

Ex. of Ionic Bonds: Determine charge & CRISS CROSS

Na Cl

Al Br

K O

Mg N

K P

Na+Cl- = NaCl

Al3+Br- = AlBr3

K+O2- = K2O

Mg2+N3- = Mg3N2

K+P3- = K3P

DETERMINE THE CHARGE OF THE ION, CRISS CROSS CHARGES so the compound is NEUTRAL!!! (compounds DO NOT have charges)

Cation + always goes 1st, Anion – always goes last!!

Making Ionic Compounds

+ CATION always goes first in a compound!!!

- ANION always goes last in a compound!!!

Another Example: reducing to the simplest formula

What is the ionic compound formula for the calcium ion with the oxide ion?

Ca O

Ionic Compounds w/ polyatomic ions

POLYATOMIC:DO NOT CHANGE THE FORMULA OF A POLYATOMIC ION!!!

Determine the formula for:

Li+ and PO43-

NH4+ and N3-

More practice

Determine the charge on an Aluminum ion, then pair it with the sulfite ion, SO3

2-

More practice…

Determine the charge on a Calcium ion, then pair it with the sulfate ion, SO4

2-

Covalent Bonds

The joy of sharing!

Covalent Bonds

Covalent bonds: occur between two or more nonmetals; electrons are shared not transferred (as in ionic bonds)

The result of sharing electrons is that atoms attain a more stable electron configuration.

Covalent Bonds

Most covalent bonds involve: 2 electrons (single covalent bond), 4 electrons (double covalent bond, or 6 electrons (triple covalent bond).

MOLECULES

Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)

Octet Rule

Octet Rule: The representative elements achieve noble gas configurations (8 electrons) by sharing electrons.

THERE ARE A FEW EXCEPTIONS!!!Hydrogen can only have 1 bond (2 electrons around it)


H2 HBr

CCl4 O2

N2 CO


H2 HBr

CCl4 O2

N2 CO


H2 HBr

CCl4 O2

N2 CO


H2 HBr

CCl4 O2

N2 CO


H2 HBr

CCl4 O2

N2 CO


H2 HBr

CCl4 O2

N2 CO

Writing Lewis Formulas (for molecules) How to:

1. Add up all valence electrons for EACH atom in the molecule

2. Attach atoms with a single bond (skeleton drawing) (C = ALWAYS CENTRAL, H = ALWAYS ON OUTSIDE OF STRUCTURE)

3. Subtract out 2 electrons for each single bond you drew (EACH BOND = 2 electrons)

4. Distribute remaining electrons (in pairs) around atoms to obtain octet rule (except H)

5. If there’s not enough electrons to satisfy the octet rule, make MULTIPLE BONDS (double, triple)

Now let’s get more complex…

Lewis Structure Examples (remember your 5 steps):

CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+


CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+


CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+


CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+


CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+


CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+


CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+


CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+

Electronegativity

We’ve learned how valence electrons are shared to form covalent bonds between elements. So far, we have considered the electrons to be shared equally. However, in most cases, electrons are not shared equally because of a property called electronegativity.

Electronegativity

The ELECTRONEGATIVITY of an element is: the tendency for an atom to attract electrons to itself when it is chemically combined with another element.

The result: a “tug-of-war” between the nuclei of the atoms.

Electronegativity

Electronegativities are given numerical values (the most electronegative element has the highest value; the least electronegative element has the lowest value)

*** see Periodic Table, in lower part of element box

Most electronegative element:Fluorine (4.0)

Least electronegative elements: Fr (0.70), Cs (0.79)

Electronegativity Notice the periodic trend:

As we move from left to right across a row, electronegativity increases (metals have low values nonmetals have high values – excluding noble gases)

As we move down a column, electronegativity decreases.

The higher the electronegativity value, the greater the ability to attract electrons to itself.

Nonpolar Bonds

When the atoms in a molecule are the same, the bonding electrons are shared equally.

Result: a nonpolar covalent bond Examples: O2, F2, H2, N2, Cl2

Polar Bonds When 2 different atoms are joined by a

covalent bond, and the bonding electrons are shared unequally, the bond is a polar covalent bond, or POLAR BOND.

The atom with the stronger electron attraction (the more electronegative element) acquires a slightly negative charge.

The less electronegative atom acquires a slightly positive charge.

Polar Bonds

Example: HCl

Electronegativities: H = 2.20 Cl = 3.16

H Cl

-+

Polar Bonds

Example: H2O

Electronegativities: H = 2.20 O = 3.44

Polar Bonds

Example: H2O


Polar Bonds

Example: H2O


Predicting Bond Types Electronegativities help us predict

the type of bond:

Electronegativity Difference

Type of Bond Example

0.00 – 0.40 H-H

0.41 – 1.00 H-Cl

1.01 – 2.00 H-F

2.01 or higher Na+Cl-

covalent

(nonpolar)covalent

(slightly polar)covalent

(very polar)

ionic

Polar Molecules A polar bond in a molecule can

make the entire molecule polar

A molecule that has 2 poles (charged regions), like H-Cl, is called a dipolar molecule, or dipole.

Polar Molecules The effect of polar bonds on the polarity of

a molecule depends on the shape of the molecule.

Example: CO2

O = C = O shape: linear

*The bond polarities cancel because they are in opposite directions; CO2 is a nonpolar molecule.

Polar Molecules The effect of polar bonds on the

polarity of a molecule depends on the shape of the molecule.

Water, H2O, also has 2 polar bonds: But, the molecule is bent, so the bonds

do not cancel. H2O is a polar molecule.

RESONANCE STRUCTURES

Resonance

A molecule or polyatomic ion for which 2 or more dot formulas with the same arrangement of atoms can be drawn is said to exhibit RESONANCE.

Resonance Example CO3

2-

3 resonance structures can be drawn for CO32-

the relationship among them is indicated by the double arrow.

the true structure is an average of the 3.


2-



the true structure is an average of the 3.


2-



the true structure is an average of the 3

Resonance Structures

Another way to represent this is by delocalization of bonding electrons:

(the dashed lines indicate the 4 pairs of bonding electrons are equally distributed among 3 C-O bonds; unshared electron pairs are not shown)

VSEPR

valence shell electron pair repulsion

Molecular Shape Lewis structures (electron dot

structures) show the structure of molecules…but only in 2 dimensions (flat).

BUT, molecules are 3 dimensional! for example, CH4 is:

Molecular Shape Lewis structures (electron dot structures)

show the structure of molecules…but only in 2 dimensions (flat).

BUT, molecules are 3 dimensional! but in 3D it is:

a tetrahedron!= coming out of

page= going into page= flat on page

Why do molecules take on 3D shapes instead of being flat?

Valence Shell Electron Pair Repulsion theory

“because electron pairs repel one another, molecules adjust their shapes so that the valence electron pairs are as far apart from another as possible.”

Why do molecules take on 3D shapes instead of being flat?

Valence Shell Electron Pair Repulsion theory

Remember: both shared and unshared electron pairs will repel one another (unshared electron pairs repel MORE than shared electron pairs in bonds)

H—N — H—

H

Non-BondingPairs

BondingPairs

5 Basic Molecule Shapes

1. tetrahedral

example: CH4


2. Pyramidal

Example: NH3

(note: unshared pair of electron repels, but is not considered part of overall shape; no atom there to contribute to the shape)


3. Bent or angular

Example: H2O Notice electron pair repulsion


4. Linear

Example: CO2


5. Trigonal planar or planar triangular

Example: BF3

CHEAT SHEET…Let’s make this easy to remember shapesLinear= only 2 regions of SHARED electron pairs

(no unshared electrons) around the CENTRAL atom

Bent = 2 SHARED electron pairs, 2 UNSHARED electron pairs around the CENTRAL atom

Tetrahedral = 4 SHARED electron pairs (meaning 4 bonds) around the CENTRAL atom

Trigonal Planar = 3 SHARED electron pairs (meaning 3 bonds) around the CENTRAL atom

Pyramidal = 3 SHARED electron pairs, 1 UNSHARED pair of electrons around the CENTRAL atom

Geometry and polarity

Three shapes will cancel out polarity. Linear = CO2 = NONPOLAR


Three shapes will cancel out polarity. Trigonal Planar = BF3 = NONPOLAR

120º


Three shapes will cancel out polarity.Tetrahedral = CH4

= NONPOLAR


Others don’t cancel Bent = H2O = POLAR


Others don’t cancel Pyramidal = NH3 = POLAR

BONDING General Rule of Thumb: metal + nonmetal = ionic polyatomic ion + metal or polyatomic ion = ionic (both) nonmetal + nonmetal(s) = covalent.

Documents

number of electrons

octet rule atoms of

number of valence electrons

octet rule noble gas

elements atoms

elements na

ne atoms of nonmetals

pseudo noblegas slide