1. The Structure of Matter
How Matter Combines and What It May Make…

2. Compounds And Mixtures
What are compounds?
Two or more elements bonded together
Chemical bonds distinguish compounds from mixtures
The force that holds two atoms together
The ability to write a formula, such as H2O, indicates a compound
A compound always has the same chemical formula

3. Types of Bonds Ionic: Opposing ions attract one another
Covalent: Atoms share e-
Metallic bonds: Metal atoms bonding with other metal atoms
Hydrogen: Bonding via hydrogen or polar molecules
Bonds can bend and stretch without breaking (bond lengths are averages)

4. Models of Compounds Ball and Stick Space-filling Structural formula
good for showing angles
Space-filling
good for relative atomic sizes
Structural formula
most commonly used
Caffeine
Various Hydrocarbons

5. How Does Molecular Structure Affect Properties?
Recall that ionization energy refers to how easily an atom loses an electron
Electron affinity refers to how much attraction an atom has for electrons
Nobel gases have a full valence shell
Thus, they have high ionization energy and low electron affinity – THEY ARE INERT!!!
Means they don’t react or don’t want to change
Elements tend to react to acquire the stable electron structure of a noble gas

6. Formation of Ions Cation Anion
By losing an electron, the atom has gained a positive charge (+1) and now has a different electron configuration
Anion
Elements on the right side of table have high electron affinity and will therefore accept electrons forming a negative charge (-1)

7. Ionic Bonds
A type of chemical bond that holds oppositely charged particles together in an ionic compound
Oxides: Metal and Oxygen (O)
Salts: Metal and non-metals
Electrons are transferred, not shared
Ionic compounds are in the form of “networks”, not “molecules”

8. Ionic Bonds
During the formation of ionic bonds, the positive and negative ions are packed into a regular repeating pattern that balances the forces
This strong + to – attraction in ionic bonds create a crystal lattice
This causes all the cations to be completely surrounded by anions and vice versa

9. Properties of Ionic Compounds
Crystalline solids – STRONG BONDS
Hard and brittle
High melting points
High boiling points
High heats of vaporization
High heats of fusion
Good conductors of electricity when molten - dissolved
Poor conductors of heat and electricity when solid
Many are soluble in water

10. Energy of the Ionic Bond
Creating ionic bonds is always exothermic
“exo” means out or exit, “thermic” means thermal or heat
Heat or energy is released
If equal amount of energy were inserted into the bond, then the ionic bond would break
The more negative the lattice energy, the stronger the force of attraction
Larger absolute charges and smaller ions have a larger negative lattice energy

11. Covalent Bonds
A chemical bond that results from sharing of valence electrons
The shared electrons are considered to be part of the complete energy level of both atoms
This is also known as a molecule
These tend to form between nonmetals

12. Covalent Properties
Covalent bond involves attractive and repulsive forces
With molecules, it is endothermic to break bond and exothermic when building it, same as ionic
Covalent bonds can form two different types: Polar and non-polar

13. Covalent Bonds
Example of these can be found in the naturally occurring diatomic molecules
Remember… Mr. BrINClHOF
Bromide (Br), Iodine (I), Nitrogen (N), Chlorine (Cl), Hydrogen (H), Oxygen (O) and Fluorine (F)
Single Covalent Bonds
This occurs when a single pair of electrons is shared
The shared electron pair, or bonding pair, is represented by either a pair of dots or a line
Lewis Dot Structures

14. Multiple Covalent Bonds
Atoms can attain a noble-gas configuration by sharing more than one pair of electrons between two atoms
Mostly formed by C, N, O and S
This can be seen with double and triple covalent bonds

15. Chemical Bonds
Non-Polar Covalent bonds are between two identical nonmetal atoms
Polar Covalent bonds are between two different nonmetal atoms
Ionic bonds primarily form between nonmetals and reactive metals

16. Polar vs. Non-polar
Recall: Electronegativity is the attraction of electrons via atoms bonded together in molecules
If the difference in electronegativity between the bonded atoms is great enough, then the electrons will orbit more frequently around the atom with greater electronegativity
Equal sharing of electrons are non-polar molecules
Unequal sharing creates polar molecules

17. Polar Molecules
Polar molecules will have the uneven distribution of electrons because of the nucleus’s pull
This creates POLES to the molecule similar to the Earth’s North and South Poles
e.g.: HF – Hydrogen (2.2) vs Fluorine (4.0)
Electrons will be found closer to fluorine than hydrogen, creating a negative pole are the fluorine nucleus
Thus, a positive pole near hydrogen
This created a DIPOLE molecule (two poles)
+ -

18. Covalent vs. Ionic Different Different Alike Topic Topic Share
electrons
(polar vs. nonpolar)
Transfer
electrons
(ions formed)
+ / -
Alike
Chemical
Bonds
Topic
Topic
Between
Two
Nonmetals
Between
Metal and
Nonmetal
Covalent
Ionic
Electrons
are
involved
Weak
Bonds
(low melting point)
Strong
Bonds
(high melting point)

19. Metallic Bonds Electrons move freely between metal atoms
An attraction between nuclei and neighboring electrons packs the metals in tightly, and allows electrons to move freely between
This is why metals conduct electricity

20. Metallic Bonding Cations “electron sea” e1- + Free electrons
Metallic bonding is the attraction between positive ions and
surrounding freely mobile electrons
Most metals contribute more than one mobile electron per atom

21. Shattering an Ionic Crystal; Bending a Metal
broken crystal
An ionic crystal
Force + - + -
Electrostatic
forces
of repulsion + - + -
A metal
Force + +
No electrostatic forces of repulsion –
metal is deformed (malleable)

22. Chemical Bonds (The Glue) Ionic Bonds
What holds ionic bonds together?
The positive cation(s) and the negative anion(s) attraction via the electromagnetic force
IMPORTANT: There are four universal forces
Strong Force: holds nucleus together (quarks or nucleon glue)
Electromagnetic Force: attracts opposite charges and poles
Weak Force: allows transmutation of nucleons
Gravity: mass attracts mass

23. Chemical Bonds (The Glue) Covalent Bonds
Van der Waal’s Forces
What holds a polar covalent molecule together?
The partial positive and partial negative charges via the electromagnetic force (again)
Also known as dipole forces
Then what holds a non-polar covalent molecule together?
London Dispersion Forces
Electrons are always moving! Constantly moving within their own orbits and around the nuclei of the atoms in the bonds
At any point, there will be a congregation of electrons in one area more than another
This creates small TEMPORARY dipole forces that last fractions of fractions of a second

24. Hydrogen Bonding
This type occurs when hydrogen has a partial positive charge from the unequal distribution of electrons
Electronegativity
The partial positive hydrogen will bind with a partial negative atom on another molecule
This connects water to water for example
Dipole-dipole (two polar molecules binding)

25. Polyatomic Ions Within Compounds
Groups of covalently bonded atoms that have either lost or gained electrons
Behave as ions (cations or anions)
Parentheses group the atoms of a polyatomic ion and show that they act as any elemental ion
Cations are placed in front and anions behind
The charge of the polyatomic ion is for the entire group, not just part of it
Names may be related to oxygen content

26. Examples of common polyatomic ions:
1- charge
Formula
Name
C2H3O21-
acetate
ClO31-
chlorate
ClO21-
chlorite
CN1-
cyanide
HSO31-
hydrogen sulfite
HSO41-
hydrogen sulfate
HCO31-
hydrogen carbonate
OH1-
hydroxide
NO31-
nitrate
NO21-
nitrite
2- Charge
Formula
Name
CO32-
carbonate
CrO42-
chromate
SO42-
sulfate
SO32-
sulfite
3- charge
Formula
Name
AsO43-
arsenate
C6H5O73-
citrate
PO43-
phosphate
PO33-
1+ charge
Formula
Name
NH41+
ammonium

27. Compounds with Polyatomic Ions
Ba(CO3)2
NH4Cl
Ba(OH)2
K(ClO3)
Fe(OH)3
Fe(OH)2
Al(OH)3
Al2(SO4)3
Notice the position of the subscripts
If within the ( ) then they are a part of the polyatomic ion, if not, then it is stating the number of the polyatomic ions
Al2(SO4)3
3 polyatomic ions
Total of 2 Al’s, 3 S’s, and 12 O’s

28. Compounds with Polyatomic Ions
Ba (OH)2
The metal (Ba) is held to the polyatomic ion by ionic bonding
These nonmetals are held to each other covalently

29. Naming Ionic Compounds:
Charge is determined mathematically
Overall charge must be zero unless otherwise stated
Fe2O3 has iron and oxygen
Cation is first and anion is second
If there isn’t a polyatomic ion, then it ends in –ide
All cations need a roman numeral unless its charge is constant
i.e. transition metals can have changing charges
Each O is -2 (periodic table)…
O’s equal a total of -6, so Fe’s must equal +6…
There are 2 Fe’s, so each must be +3…
Fe2O3 is “Iron III Oxide”

30. Examples Fe2O Answer: Iron I Oxide W2O 7 --- W2 O7
Answer: Tungsten VII Oxide
Hg(CN)2 --- Hg (CN)2
Answer: Mercury II Cyanide
First, -14 this side
Third, 2 W’s must total +14, so each one is +7
Second, +14 this side
First, -2 this side from the polyatomic ion chart
Third, 1 Hg must be +2 by itself
Second, +2 this side

31. Naming Covalent Compounds
Numerical prefixes are used to name covalent compounds of two or more elements
Mono- , Di- , Tri- , Tetra- , Penta- , Hexa- , Hepta- , Octa- , Nona- , Deca-
Elements are listed left to right on the periodic table and ends in –ide
Silicon dioxide --- SiO2
If the first element is singular, “mono” may be left off
Dihydrogen Oxide --- H2O (water)
Nitrogen Trichloride --- NCl3

32. Naming Basic Acids Look for two characteristics:
The compound/molecule will begin with hydrogen, H-
This will identify it as an acid
Look for what is immediately after the hydrogen,
If it is a single element, then it is a binary acid
If it is a polyatomic ion, then it is an oxyacid
Binary Acids: HCl, HBr, HI, etc.
Called “hydro-” name of the anion with an –ic ending, followed by acid
Hydrochloric acid, hydrobromic acid, hydroiodic acid, etc.

33. Naming Basic Acids
An oxyacid is when there is a polyatomic ion after the hydrogen, H2SO4
If the polyatomic ion ends in –ite, then the acid will end in –ous
If the polyatomic ion ends in –ate, then the acid will end in a –ic ending
HNO3 – Nitric Acid
HNO2 – Nitrous acid

34. VSEPR
You should now be familiar with bond types, naming and drawing them in 2-D (Lewis Dot Structures)
Now we will begin to picture them in 3-D
3-D models are based on VSEPR:
Valence Shell Electron Pair Repulsion Theory

35. VSEPR
The idea behind VSEPR is that covalent bonds and lone pair of electrons spread to the furthest distance they can from one another
Covalent bonds consist of electrons and they have the same charge
Thus, VSEPR explains why molecules have their shapes

36. VSEPR Flow Chart

37. VSEPR Flow Chart Draw the Lewis Structure for the compound in question
Count the number of “things”
Follow the flow-chart

38. Molecule Shapes Tetrahedral Trigonal pyramidal Trigonal planar
Tetrahedral molecules look like pyramids with four faces and each point corresponds to an atom that's attached to the central atom
Bond angles are degrees
Trigonal pyramidal
It's like a tetrahedral molecule, except flatter
Bond angles are degrees (it's less than tetrahedral molecules because the lone pair shoves the other atoms closer to each other)
Trigonal planar
It looks like the hood ornament of a Mercedes automobile
The bond angles are 120 degrees

39. Molecule Shapes Bent Linear They look bent
Bond angles can be either 118 degrees for molecules with one lone pair or degrees for molecules with two lone pairs
Linear
The atoms in the molecule are in a straight line
This can be either because there are only two atoms in the molecule (in which case there is no bond angle, as there need to be three atoms to get a bond angle) or because the three atoms are lined up in a straight line (corresponding to a 180 degree bond angle)