unit 9 bonding. the attractive forces between atoms leads to chemical bonds that result in chemical...

Unit 9 Bonding

The attractive forces between atoms leads to chemical bonds that result in chemical compounds.

CH4 methane gas molecule

Why do atoms form bonds?

Atoms form bonds due to the need to have the most stable configuration for its electrons. Atoms lose, gain, or share valence electrons in order to achieve a lower energy state (stable).

How atoms bond with each other depends on:

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ElectronegativityIonization Energy# Valence Electrons

Metallic bonding

http://www.launc.tased.edu.au/online/sciences/PhysSci/pschem/metals/Metals.htm

Metallic bonding

Metallic bonding is the strong attraction between closely packed positive metal ions and a 'sea' of delocalized electrons.

http://www.bbc.co.uk/schools/gcsebitesize

What are the properties of metals?

How does metallic bonding affect the properties of metals?

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What property of metals is illustrated above?

Metallic bonds

Ionic Bonding

• an electrostatic force – Electrostatic refers to the attraction between

opposite charges

• Stronger than metallic bonds because of the opposite charges

Ionic bond

• Formed by the transfer of electrons between a metal and a nonmetal

• # of e- lost by metal = # of e- gained by nonmetal

Charge

• The number of electrons that need to be lost or gained by an atom so it has the same electronic configuration as a noble gas.

• When an atom gains or loses electrons, it is called an ION.– + ions are cations– - ions are anions

• The charge of the ion is called the OXIDATION NUMBER.

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Charge based on “copy cat” principleSulfur wants to be like Argon (1s22s22p63s23p6)

Sulfur Atom 1s22s22p63s23p4

16 Protons 16(+)16 Electrons 16 (-)

0 No charge

Sulfur has 6 valence e- and will gain 2 more to complete an octet. The result is a:Sulfur Anion 1s22s22p63s23p6

16 Protons 16 (+)18 Electrons 18 (-)

2- Charge

Ionic Bonds

Strong electrostatic (positive-negative) force in ionic compounds makes a strong ionic bond.

How does the strong ionic bond affect the properties of ionic compounds?

Formula unit – smallest unit of an ionic compound; lowest whole number ratio of ions represented in an ionic compound

• The greater the difference between eN values of 2 atoms is, the more ionic the bond will be. We call this “ionic character”.

0.8- 4.0 = 3.2 very ionic

Ionic Bond

http://www.chm.bris.ac.uk/pt/harvey/gcse/ionic.htm

Na+ Cl-

Writing Lewis Dot Structures

• Element symbol represents the kernel (core) of the atom (nucleus and inner e-)

• Dots represent the valence e-

www.meta-synthesis.com 

• Metals tend to lose e- while nonmetals tend to gain electrons

Writing Lewis Dot Structures - Ionic

hyperphysics.phy-astr.gsu.eduIonic bonds

• Metals tend to lose e- while nonmetals tend to gain electrons

Writing Lewis Dot Structures – Ionic Bonds

chemistry58.wikispaces.com

Ionic bonds

Properties of ionic compounds (all related to the strong attraction between

the + and – charges)• high melting points and

boiling points• hard solids• good conductors – in

aqueous solutions and when molten

• have a crystal lattice structure Ions are

here

Really, we don’t hate you.

Covalent Bond

• Covalent Bond –formed when two nonmetals share pairs of valence electrons in order to obtain the electron configuration of a noble gas

• Molecule - formed when two or more atoms bond covalently. –(A molecule is to a covalent bond as a formula unit is to an ionic bond.)

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How covalent atoms bond

Diatomic Molecules• HOFBrINCl

Share electrons when they bond together23

Polyatomic Ions

• covalently bonded group of atoms, with a charge

• Watch this

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They are listed on your STAAR chart. You will not have a bad time.

Properties of Covalent Molecules

• Can exist as gases, liquids, or solids depending on molecular mass and polarity

• Usually have lower MP and BP than ionic compounds of the same mass

• Do not usually dissociate (break apart into ions) in water

• Do not conduct electricity

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How to draw Lewis dot structures for covalent molecules.

1. Write the formula for the compound.2. Count the total number of valence electrons. 3. Predict the location of the atoms:

a) Hydrogen is NEVER the central atom.b) If carbon is present, it is ALWAYS the central atom.c) If there is only 1 atom of an element, it is the central atom.d) The least electronegative atom is generally the central atom.

4. Place one electron PAIR between the central atom and each ligand (side atom) to “hook” the atoms together.

5. Dot the remaining electrons in pairs around the compound to complete the octet. Start with the ligands.

6. Check that each atom has an octet. (H only needs a pair, not an octet.)7. Watch This

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Lewis Structures for Molecules

• Draw the Lewis dot structure for these molecules:– Hydrogen + Bromine (HBr)

– Carbon + Chlorine (CCl4)

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Bonding e- PairsLone Pairs(nonbonding electrons)

Covalent bonds

Writing Lewis Dot Structures - Covalent Bonds

Number of bonds• Single Bonds - when one pair of e- is shared

between atoms

• Double bond – when atoms share 2 pairs of valence electrons; ex. O2

• Triple bond – when atoms share 3 pairs of valence electrons; ex. N2

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Describing bonds

• Sigma bond - the first bond between 2 atoms– A single bond is a sigma bond.

• Pi bond - the second bond between 2 atoms– A double bond consists of a sigma bond and a pi

bond. – A triple bond consists of a sigma bond and two pi

bonds.

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Carbon can form single, double and triple bonds with itself.

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• The shape of a molecule plays a very important role in determining its properties.

Why are molecular shapes important?

• Properties such as smell, taste, and proper targeting (of drugs) are all the result of molecular shape.

VSEPR Theory also called electron geometry

• Electron groups around the central atom will be most stable when they are as far apart as possible. We call this valence shell electron pair repulsion theory.– Because electrons are negatively charged, they

should be most stable when they are separated as much as possible.

• The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule.

TO DETERMINE VSEPR SHAPE (electron geometry)

1) Draw the Lewis dot structure for the molecule2) Identify the central atom3) Count the number of electron groups around the

central atom.4) Look up the VSEPR shape on the chart.**shapes with no lone pairs are symmetrical**shapes with lone pairs are assymmetrical

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O N O ••••

••

••

•••• There are three electron groups on N•One lone pair•One single bond•One double bond

Electron Groups• Each lone pair of electrons = one electron

group on a central atom.• Each bond = one electron group on a central

atom (regardless of whether it is single, double, or triple.)

Two Electron Groups: Linear Electron Geometry

• When there are two electron groups around the central atom, they will occupy positions on opposite sides of the central atom.

• This results in the electron groups taking a linear shape.

Linear Geometry

Three Electron Groups: Trigonal Planar Electron Geometry

• When there are three electron groups around the central atom, they will occupy positions in the shape of a triangle around the central atom.

• This results in the electron groups taking a trigonal planar shape.

Trigonal Planar Geometry

Four Electron Groups: Tetrahedral Electron Geometry

• When there are four electron groups around the central atom, they will occupy positions in the shape of a tetrahedron around the central atom.

• This results in the electron groups taking a tetrahedral shape.

Tetrahedral Geometry

Molecular Geometry• The actual geometry of the molecule may be

different from the VSEPR shape.• Lone pairs repel bonded atoms which distorts the

expected shape.

Bond Angle Distortion from Lone Pairs

Electron Geometry: Tetrahedral Tetrahedral TetrahedralMolecular Geometry: Tetrahedral Trigonal Pyramidal Bent

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Predicting the Shapes around Central Atoms

1. Draw the Lewis structure.2. Determine the number of electron groups

around the central atom.3. Classify each electron group as a bonding or

lone pair, and count each type.– Remember, multiple bonds count as one group.

4. Look it up

Practice:

• Determine the shape.

1. NF3

2. SiCl4

3. H2O

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Types of Bonds:• Nonpolar covalent equal sharing of

electrons between atoms; occurs between the atoms in a diatomic molecule (HOFBrINCl) and between C and H; ex. CH4

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Polar Covalentunequal sharing of electrons between atoms; occurs between two nonmetals or a nonmetal and a metalloid; ex. H2O

Electrons

Ionic

• complete transfer of electrons

• occurs between– m/nm (ex. NaCl)– m/PAI– PAI/nm– PAI/PAI

THIS IS A CONTINUUM. IT DESCRIBES THE “IONIC CHARACTER” OF THE BOND.

Bond type

Non-Polar Covalent Polar Covalent Ionic NPC PC I

• Difference in electronegativity values

• Distance between atoms on the periodic table

Small medium big

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Practice:

• What type of bond exists in each of the following?

1. HCl2. CaO

3. H2O

4.Br2

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Water is a POLAR molecule

The more electronegative atom will have a slight negative charge, the area around the least electronegative atom will have a slight positive charge. 52

Symmetric molecules tend to be nonpolarAsymmetric molecules with polar bonds are polar

53Symmetric means there are no lone pair around the central atom.

Type of Compound

Elements involve in bonding

(metal/non metal)

Valence electrons

are…

Melting /Boiling

point

Electrical conductivity

Other properties

Metallic

Ionic

Covalent

Type of Compound

Elements involve in bonding

(metal/non metal)

Valence electrons

are…

Melting /Boiling

point

Electrical conductivity

Other properties

Metallic Metal-metal delocalized high Conductor Malleable, ductile, shiny

Ionic Metal - nonmetal

Lost/gained high Conducts in solutions or molten

Brittle, solid at room temperature

Covalent Nonmetal - nonmetal

shared low Non-conductor

Mostly liquid or gas at room temperature

Bonding determines some physical properties