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Electron dot symbols

Electron Dot Symbols

Hydrogen would look like: H½ Helium would look like: He: Helium breaks the rule because it has a filled outer shell. Carbon would look like: . C .

Structure of Molecules

Bonding and Shape

Electron dot symbols are a way of showing the valence electrons for bonding Valence electrons are the electrons in the outermost energy level. To draw the electron dot symbol, we imagine a square around the chemical symbol and place the valence electrons around the square. One electron per side.

. .

Electron Dot Structures

Also called Lewis Structures

Electrons and ions

Metals lose electrons

Charge will equal group number for main group metals (tall columns)

How atoms connect together to make compounds. Two types:

Ionic compounds Covalent compounds

Non-metals gain electrons

Charge will equal group number minus eight

We just saw that the atoms will gain or lose electrons so that it has an electron configuration that looks like a noble gas. How does this affect the electron dot symbol?

Structure can give us other properties of the compound

Charge is the number of electrons gained or lost.

Lithium:

Li

Fluorine:

F

F

Lithium Ion: Li

+

Fluoride Ion:

1

Electron dot structure of ionic compounds

Write the symbols for the ions. Place them next to each other

Polyatomic ions

A group of atoms, usually non-metals, that are covalently bonded together but have an overall charge. The Lewis structure takes this charge into account

Cations (positive charge) have fewer electrons Anions (negative charge) have more electrons

Naming ionic compounds

The name of ionic compounds consists of two parts, the name of the cation and the name of the anion. The cation is always named first.

Cation name Anion Name

Li

+

F

Example: NaCl Sodium Chloride

Naming ionic compounds

The names of metal ions are same as the metals. But there are two kinds of metal ions: Type I and Type II

Type I metals form only one kind of ion. These are: Groups IA and IIA, Al, Ga, Zn, Cd, and Ag. Their charges are equal to their group number. Type II metals form more than one kind of ion. For our purposes, if it's not Type I it's Type II. Their name contains a Roman numeral to specify the charge on the ion.

Naming ionic compounds

The names of non-metal ions consist of the root of the non-metal name with the ­ide ending.

O2- is Oxide, N3- is Nitride

The name of the compound gives us the formula. For example: Lithium Oxide We know that there is a lithium ion (Li+) and an oxide ion (O2-) The charges give us the ration of the ions in the compound. How is that? We can use the crossover method: Li+ O2The charges are crossed-over and this gives us the formula Li2O. "1"s are never explicitly written.

The name of the compound gives us the formula. For example: Lithium Oxide

2

More on Type II metals

Iron is a type two metal; it forms Fe2+ and Fe3+ ions. The Fe2+ is called Iron(II) ion. It used to be called Ferrous ion. The newer naming system gives us the charge without having to memorize it. The Fe3+ is called Iron(III) ion and used to be called Ferric ion. The thing to remember is that the number in parenthesis is the charge, NOT the number of that kind of atom in the compound!

How do we get the charges on Type II metals?

Reverse the cross-over method. Example: Mn3P2

Polyatomic ions

List on handout must be memorized. (Names, formulas and charges) Most of them are anions, only one is a cation that appears in compounds. The other cation is only in aqueous solution. A couple contain metals

Mn3P2 We know that Phosphide has a 3- charge, it's in group VA. This means that the Manganese ion must have a 2+ charge. This gives us the name Manganese(II) Phosphide. If there are no subscripts, the charges on both have equal magnitude but opposite sign.

Naming ionic compounds (reprise)

Naming ionic compounds that contain polyatomic ions is no different than naming regular ionic compounds. Example: Fe(C2H3O2)2

Iron(II) Acetate (or Ferrous acetate)

Electron dot structures for covalent compounds

Determine the skeletal structure of the compound. Count the number of valence electrons in the compound. Connect all the outer atoms to the central atom by single bonds (use a single line). Give all of the outer atoms an octet by placing the appropriate number of electron pairs around each atom. Any left over electron are placed in pairs on the central atom. If the central atom does not have an octet, form double or triple bonds until the central atom has an octet.

Example

CCl4

3

Determine skeletal structure

Atom that is able to form the most number of bonds is the central atom. In this case, carbon.

Count valence electrons

Carbon has 4 Each chlorine has 7 (7x4=28) Total is 32 (28+4=32)

Connect outer atoms with single bonds

Each single bond is two electrons. There are a total of 8 electrons used here.

Cl Cl C Cl Cl

Cl Cl C Cl Cl

Fill octets on outer atoms

Each of the chlorines needs six more electrons for an octet.

Check structure for octets

All the atoms have an octet so this is the completed structure. No multiple bonds are needed.

Naming Covalent compounds

Not as involved. Names are more descriptive of the formula Must memorize the list of prefixes Needed to distinguish between the multitude of compounds between different non-metals We will only name binary covalent compounds (those containing only 2 elements)

Cl Cl C Cl Cl

Cl Cl C Cl Cl

This structure uses all of the available electrons.

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Prefixes

1: mono 2: di 3: tri 4: tetra 5: penta 6: hexa 7: hepta 8: octa 9: nona 10: deca

Rules for using prefixes

Mono prefix is never used with the first element. If the prefix ends with the same sound as the beginning of the element, the end of the prefix is dropped. Like ionic compounds, the second element is named with the root of the element name with the appropriate prefix and the ­ide suffix.

Examples

NO2

Nitrogen Dioxide

SO3

Sulfur Trioxide

P2O5

Diphosphorus Pentoxide

Exceptions

H2O is water not Dihydrogen Monoxide NH3 is Ammonia not Nitrogen Trihydride Compounds containing mainly Carbon and Hydrogen are named under separate rules.

Shapes of molecules

Shape depends on the arrangement of electrons around the central atoms. The electrons will arrange themselves so that they are as far from each other as possible. This arrangement of electrons is what we perceive as the shape of the molecule.

VSEPR

Count the number of groups of electrons around the central atom Determine the number of lone pairs out of that number of groups Follow the table (which must be memorized!).

5

VSEPR

Number of groups 2 Number of lone pairs 0 1 0 1 0 4 1 2 Tetrahedral EPG Linear MG Linear Linear Trig. Planar Bent Tetrahedral Trig. Pyramidal bent

Electronegativity

Electronegativity is the tendency of an element to pull electrons when it's in a bond with another element. Fluorine is the most electronegative, Cesium is the least. Electronegativity decreases as you move away from Fluorine in the periodic table The three most electronegative elements are N, O and F (in increasing order).

Molecular Polarity

Some molecules have an uneven distribution of electrons due to a difference in electronegativity between the elements present in the compound. If the molecule is asymmetric, it is polar. For this class, if it contains a lone pair it is polar. If the molecule is symmetric, it's nonpolar.

3

Trigonal Planar

What do we mean by symmetry?

H H C H H

Methane (CH4) is highly symmetric. No matter how we look at it, it looks the same. Bromomethane (CH3Br) is asymmetric because it looks different from different directions.

H H C Br H

Rotate by 120°

H H C H Br

H H C Br H

Because this molecule is asymmetric, it is polar. There is an uneven distribution of electrons in the molecule (they spend most of the time around the bromine).

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