Chemical Bond: mutual electrical attraction between the ______ & ____________ of different atoms that bond together. The type of bonding is determined by the way the valence e-’s are ____________. Ch.6 Chemical Bonding

Two Types of Bonds Ionic: results from 1 atom giving up its valence e-’s (  cation) & __________ them to another atom (  anion)e.g. Covalent: results from the ______ of valence e-’s between 2 atoms Most bonds are between these extremes!

Ionic bond Complete transfer of an electron making one atom ______ (Cation) and one _______ (Anion) They then attract each other like magnets because of their opposite _______ Only one has electron density

Non-Polar Covalent Non-Polar Covalent: the bonding valence e-’s are _______ _____ by the atoms resulting in equal distribution of electrical charge. e.g.

Polar covalent bond Polar Covalent: the bonding valence e-’s are more strongly attracted to the ____ __ ____ resulting in an _______ distribution of the valence e-’s. It is still sharing, not a transfer like in ionic. e.g.

So: Chemical Bonding The type of bonding can be determined simply by the __________ in the ______________ ( ∆EN) of the 2 atoms. (see next slide or p.177 Table 6.2 to see values)

E.g.’s: A H-F molecule has an EN difference of: (for F) – (for H) = For Na-Cl the EN difference is: (for Cl) – (for Na) = For C-O the EN difference is: (for O) – (for C) = For H-H (H 2 ) the EN difference is: (for H) – (for H) =

The difference tells you what type of bonding that is occurring: > 1.7= < 0.3= 0.3  1.7= or: EN difference = I I I I (see also p.238 Table 8.3)

Going back to the previous examples: H-F ∆ EN =  Na-Cl ∆ EN =  C-O  EN =  H-H ∆ EN =  Other examples: Mg-S ∆ EN = CO 2 ∆ EN =

Covalent Molecules Molecule: neutral group of atoms held together by ___________ bonds. e.g. Diatomic Molecule: molecules containing only 2 atoms. e.g.

More definitions… Octet Rule: where chemical compounds tend to form, such that each atom achieves an _____ of electrons in its valence shell. This is done by (becoming an ion or entering a covalent bond)

Bond Energy: the amount of energy needed to break a

e.g. Fluorine gas exists as F 2. (F 2 orbital diagram ) Bonding can ONLY occur btwn Orbitals. Each F atom has achieved a stable octet by of electrons. (e.g. NF 3 )

Electron Dot Notation: Instead of drawing the orbital diagram, which can be long & complex, there is an easier way to represent the atoms. Electron Dot Notation: is an electron notation in which only the _______________are shown, and are represented by ____________ around the element’s symbol.

Orbital Diagram to Electron Dot Diagram valence shell

Valence electrons Dot notation

Bonding When atoms bond they share their _________ electrons In dot notation this is represented as _______________________between symbols, one from each atom.

The unshared pairs of electrons are also known as ___________. The shared pairs can be represented by a _______. These representations are known as: _________________: which show the shared pairs as dots (or dashes) and the unshared pairs as dots.

The dots representing the lone pairs can also be dropped. The new representation is known as a _____________________.

A single shared pair is known as a _________ Bond. Let’s consider O 2 : The sharing of 2 pairs of electrons between 2 atoms is known as a ____________ Bond.

Let’s consider N 2 : The sharing of 3 pairs of electrons between 2 atoms is known as a _________ Bond. Double & Triple Bonds are also known as: _________ Bonds.

e.g. HCl Chlorine has a _________, but H does not. That’s because there are some exceptions due to elements just not having enough i.e. How many electrons would it need to fill an octet? Is that possible? Elements with Z<6

We still haven’t explained why carbon can form 4 bonds instead of 2… _______________ Let’s look at Carbon (6): It makes sense to assume that Carbon forms __ __________ bonds.

But when Carbon bonds with other atoms, a special thing happens. The 2s & 2p merge together to form an ________. Now apply _______ rule. So now, Carbon has ________bonds. The same applies to Si, but with the _________

Hybridisation also applies to Be & B. Beryllium forms an ________.

Boron forms an __________.

Let’s review with some examples: Draw the Lewis structure & Structural Formula for: NH 3, HCN, POI, C 2 H 6, C 2 H 4, C 2 H 2

Molecular Structure: The VSEPR Model Molecular Structure: is the of the atoms in a molecule. (i.e. the shape) There are several methods. We shall look at one simple model and deal with simple molecules only. VSEPR: helps predict the geometry (shape) of a molecule. The idea is that the valence shell electron pairs, whether bonding or lone prs, will position themselves as far apart as possible, to minimise

Molecular Shape: A 4 step program by VSEPR Step 1: Draw of the molecule Step 2: Count # electron “groups” around the central atom. Electron group = lone pr; single, double or triple bond. Step 3: Arrange electron groups as far apart as possible (to minimize ) Remember, may have to think in 3-D to get the shape of the electron arrangement. Step 4: Add “ ” atoms. Describe the shape by focusing on the atoms only. Note: Lone prs require more space than bonded prs. (see p.232-3)

e.g. BeF 2 (Note that Be only has 2 valence e-’s and no lone pairs.) Lewis Structure The 2 valence electron pairs position themselves as far apart as possible, i.e. 180 o So it has a shape/structure.

e.g. BF 3 (Note that B only has 3 valence e-’s & no lone prs.) Lewis Structure How do 3 prs of e’s position themselves as far apart as possible? This shape is known as

e.g. CH 4 But is a flat square the best shape for the valence electron pairs to be as far apart as possible? NO, we have to picture it in 3-D. This shape is known as

e.g. NH 3 N has 1 lone pr and 3 bonding prs for a total of 4 “______________.” The will position themselves in a arrangement with the 3 H’s at the end of 3 of the electron groups. The lone pr “pushes” the 3 bonding prs closer together. This shape is known as

e.g. H 2 O O has __ lone prs and __ bonding prs for a total of __ “electron groups.” The electron groups will position themselves in a __________ arrangement with the 2 H’s at the end of 2 of the electron groups. The lone pr “pushes” the 2 bonding prs closer together. This shape is known as

Summary of the various shapes derived from 4 electron groups.

Some examples: e.g. Lewis # electron 3-D Molecular Dipole structure electron arrangement structure Structure Moment groups (diagram) (shape name) CO 2 BI 3 CCl 4 PF 3 SeBr 2

Intermolecular Forces: -forces ________ molecules. -______ than ionic & covalent bonds. In Polar Covalent ( ∆ EN= ) The EN difference creates a ____ from positive to negative end. (partial charges) e.g.I-Cl=>I---Cl=>ICl (2.5) (3.0)  +  Difference Dipole

Dipole-Dipole interaction Intermolecular force based on ________ and _______ between partial charges Strong intermolecular force.

Hydrogen Bonding The strongest intermolecular force that has H partially bonded to an ______________ atom. E.g’s: Causes higher than normal boiling points  water is a ______ instead of a room temp. A type of dipole-dipole interaction but stronger e.g. I love water!!! (why?)

Non-Polar Molecules ( ∆ EN= ) There is no dipole because the EN diff is too ___. But a slight shift of the e-’s to one side creates a _________ _______, which effects the next molecule, and so on. e.g’s? London Dispersion forces (or van der Waals) Average shape Temporary shift Effects other molecules

Ionic Bonding Most of Earth’s ____ & _______ are made up of compounds held together by ___ bonds. Ionic Compound: consists of positive (___) & negative (__) ions that are combined such that the # _________ = # _________. e.g.

Most ionic compounds exist as a of alternating +’ve & -’ve ions

Formula Unit: is the simplest _____ of atoms from which an ionic compounds formula can be established. Ionic compounds can be represented by Electron Dot diagrams. e.g. NaCl Ionic crystals combine in an orderly arrangement known as a _______ _______.

Characteristics: Ionic vs Molecular -Ionic bonds are ________ than molecular (covalent) bonds Ionic compounds: -____ melting/boiling pts -very hard but _______ -when molten (or dissolved in H 2 O) they become good __________ conductors Molecular compounds: -___ melting/boiling pts -tend to be _____ at room temperature

Polyatomic Ions: -group of ________ bonded atoms that carry a ______. e.g.NH 4 + &PO 4 3- Exercise: do SO 4 ?? (what’s the charge?)

So, the Big Picture is: Bonding within molecule Between molecules

Metallic Bonding: Metals have few valence e - ’s (1,2or3) which roam freely as a “___ __ _______” throughout the metal. They do not belong to any particular atom. The atoms themselves are arranged in a _______ ____ _______. The bonding between the metal atoms & this ___ __ _______ is called metallic bonding. It is NOT bonding between different types of metal atoms.

Characteristics: Different metals can only come together when they are in their ______ _____. So now it is a mixture of 2 or more metals called an _____. E.g’s:

Alloys Substitutional Alloy: some of the ____ metal atoms are replaced by other ______ _____ metal atoms. E.g’s:Brass: Bronze: Sterling Silver: Solder: 24 Karat Gold: 18 Karat Gold:

Interstitial Alloy: some of the _____ (interstices) between the metal atoms are filled with ____ _____. E.g’s:Steel: