semester 1 review jan. 2012. significant figures counting sig figs count all numbers except: leading...

Semester 1 ReviewJan. 2012

Significant Figures Counting Sig Figs

Count all numbers EXCEPT:Leading zeros -- 0.0025Trailing zeros without

a decimal point -- 2,500

Significant Figures Calculating Multiply/Divide - The # with the fewest sig figs

determines the # of sig figs in the answer.

(13.91g/cm3)(23.3cm3) = 324.103g

324 g

4 SF 3 SF3 SF

Significant Figures Calculating Add/Subtract - The # with the lowest decimal

value determines the place of the last sig fig in the answer.

3.75 mL+ 4.1 mL 7.85 mL

224 g+ 130 g 354 g 7.9 mL 354 g

3.75 mL+ 4.1 mL 7.85 mL

224 g+ 130 g 354 g

Conversions

Change 5.6 m to millimeters

k h D d c m

starts at the base unit and move three to the right.move the decimal point three to the right

56 00

The History of the AtomBegan as the “plum Pudding” modelBasically, scientists thought that the atom was all mixed together

Ernest Rutherford RutherfordGold Foil Experimentdiscovered the nucleus!!surrounded by lots of empty space

Atomic Model of MatterBohr’s ModelProposed that electrons move in definite orbits around the nucleus

Bohr’s Model--Orbits

Atomic Model of Matter

Modern Atomic ModelAn atom has a small positively charged nucleus surrounded by a large region (electron cloud) containing enough electrons to make the atom neutral

IsotopesAll atoms of the same element have the same number of protonsHowever, sometimes they can have different numbers of neutrons Example Chlorine=17 protons, but some contain 18 neutrons and others can have 20 neutrons!These are called isotopesMost elements in the first 2 rows have isotopes

IsotopesIn nature, elements exist as a mixture of isotopesScientists use the “relative abundance” of these to calculate the average atomic mass of the elementTo indicate which isotope they are using, scientists use the complete chemical formula 1

4

Nuclear EquationNuclear equations are used to keep track of the reaction

Beginning atomic weight and atomic numberNew atomic weight and numberParticle emitted (or gamma ray)

Radioactive Half-Life In a given sample of a radioactive element, not all of the atoms will decay at the same timeDecay happens over timeThis time is called a half-life

The amount of time it takes for half of the atoms in a sample to decay

Radioactive Half-LifeCarbon-14-one of the best known radioactive isotopesHalf-life of 5730 yearsAfter 5730 years, half of the sample will have decayed (the 1st half-life)After another 5730 years, half of what was left will have decayed-you have ¼ of the original sample left. (the 2nd half-life)What will be left after the 3rd half-life?

Nuclear ReactionsScientists discovered that they could release energy

holding the nucleus together (strong force)Nuclear Fission – splitting an atomic nucleus into

two smaller nuclei of equal massUranium-235 is bombarded by a neutron and split into two separate products (krypton & barium) and additional neutronsAdditional neutrons are capable of splitting other uranium atoms creating a nuclear chain reactionNuclear chain reactions release huge amounts of energy

1. Nuclear power plants use controlled chain reactions to produce energy

2. Uncontrolled chain reactions are nuclear explosions (atomic bomb)

Nuclear FusionJust like fission, it releases a great amount of energy

Nuclear Fusion – joining two atomic nuclei to form a single nucleus

Thermonuclear reaction – requires tremendous amounts of heat (temperatures over 1 million oC)Matter is turned into plasma at these temperatures

Electronegativity (1 of 3)Electro-negativity is the ability of an atom to attract electrons that are shared in a covalent bond.

Electronegativity (2 of 3)What are the trends in electronegativity?

Atomic Radius (2 of 3)Atomic radius

trends:1) Atomic

radius increases down a group or column.

2) Atomic radius decreases across a period or row.

Ionization Energy (3 of 4)

Removing an electron becomes more difficult across a row.Removing electrons becomes easier down a column.

Electron Orbitals

Filling Orbitals

Alkali Metals• Group 1A• s orbital contains

last electron• form cations,

always 1+• soft metals• most reactive

Alkaline Earth Metals

• Group 2A• s orbital is last orbital filled• form cations, always 2+• metals• reactive

Boron Group• Group 3A• Contain one p

orbital electron.• Semimetals and

metals• All form cations• 3+ charge

Carbon Group• Group 4A• Contain 2 p orbital electrons.• C and Si generally form covalentbonds.• Ge forms a cation with 2+ charge.• Sn and Pb form 2+ and 4+ cations.• Nonmetals, metalloids, and

metals.

Nitrogen Group

• Group 5A• Contain one electron in each p

orbital.• All except for Bi can be a 3- anionor a 3+ cation.• Many can form 5+ cations.• Nonmetals, metalloids, and

metals.

Oxygen Group• Group 6A• Contain one set of paired and

two sets of unpaired p orbital electrons.

• Many form 2- anions, all but O can

form 2+ and 4+ cations.• Nonmetals, metalloids, and

metals.

Halogens• Group 7A• Halogen means “salt former.”• Contain two sets of paired p orbitalelectrons and one unpaired electron.• Commonly form diatomic molecules.F2 Cl2 Br2 I2 • Nonmetals that form anions with 1-charge.

Noble Gases• Group 8A• All s and p orbitals contain

paired electrons.• All are generally unreactive

gases.• Noble gases do not commonly

form ions.

Valence ElectronAn electron in the outermost energy level for an atom.

The electrons that interact when atoms form bonds.

H Li Na

C N F

Valence electronsThe s and p electrons in the outermost primary energy level.There are 2 valence electrons for the 1st primary energy level. (Period 1)

(there’s only an s orbital)

There are 8 valence electrons in every other primary energy level. (Periods 2-7)

(s and p orbitals)

The Octet RuleAtoms tend to gain, lose or share electrons in order to have a full set of valence electrons (usually 8)

Ionic BondingIonic Bonding – bonding that involves a transfer of electrons between atoms

One atom gains electrons (nonmetal) and the other atom loses electrons (metal)When an atom gains or loses electrons it is no longer neutral and becomes an ion

Ions – charged atomAtom that gains electrons has a negative chargeAtom that loses electrons has a positive charge

Ionic BondingFormed between cations and anionsCation=positive ionAnion=negative ionThe attraction between the positive and negative charges holds the ions together

Salt (NaCl) it’s ionic!

Ionic BondingIons in an ionic compound form a regular, repeating arrangement or crystal lattice

Chemical formula (NaCl) shows the ratio of ions present in the crystal lattice

Properties of Ionic compounds

Common properties of ionic compounds include:

High melting pointBrittlenessSoluble in waterConduct electricity when dissolvedDo not conduct when in solid form

COVALENT BONDING

Covalent bonds form between two non-metals (usually)As always, the guiding principle is the octet rule…atoms try to form an electron configuration like their nearest noble gas neighborIn this bonding scenario, both atoms share electrons to meet the octet rule (instead of transfer like ionic)

Polar Covalent Bonding

Atoms share the electron, but the sharing is unequal.

Why is the sharing unequal?One of the atoms has a higher electronegativity It pulls harder on the electron(s) than the other atomNow the electron is closer to one atom than the other

POLAR COVALENTThis results in a polar molecule-

Has ends that are differently charged (+ and -)

Overall has a neutral charge

METALLIC BONDING

Metallic bonds are made of metals (no duh!)In this type of bonding, electrons can freely move amongst metal atoms…sometimes metal bonding is called a “sea of electrons” No electron really belongs to any one atom but rather to any atom around

Metallic BondsThis sharing of electrons between all the metal atoms results in the characteristic properties of metals

Malleable-you can bend and shape it and it won’t breakGood conductors of electricityHigh melting point

*Bond Classification by Physical Characteristics…

Ionic Covalent(Polar & Non Polar)

Metallic•Crystals•Hard, Brittle•High mp•Very Soluble in Water•High Electrical Conductivity in solution

•Solids, liquids, gases.•Soft solids but many liquids and gas•Low mp•Some dissolve some don’t•Low conductivity to No Conductivity in solution

•Shiny•High Electrical Conductivity•Malleable

Rules for Lewis Structures of Molecules

1. Add up the total number of valence electrons available in the molecule.

2. The center of the molecule will be an atom that can form bonds with multiple atoms. In order of preference: C, Si, N, P, S, O.

3. Place remaining atoms around the central atom. Connect all atoms to center with single bonds.

4. Distribute remaining electrons in pairs to each atom, satisfying the octet rule (or the duet rule when appropriate).

5. Use double or triple bonds when appropriate.

Naming Ionic Compounds

The first word in the name is the name of the cation (+), and the second word is the name of the anion(-).The best way to go about naming ionic compounds is to take a look at the formula and figure out the names of the cation and anion. When you've got that, just stick them together and you've got the name of the compound.

Naming the anionIf the anion has only one atom in it, then the name of the anion is the same as the name of the elementEXCEPT the end of the element name is taken off and "-ide" is added to the end. oxygen becomes "oxide", sulfur becomes "sulfide", phosphorus is "phosphide", et cetera. If the anion has more than one atom, then we'd say that it's a "polyatomic ion", meaning that the anion has more than one atom. Look up the polyatomic ion in a table and you've got the name.Thus, OH- is "hydroxide", SO4

2- is sulfate, et cetera.

Naming Covalent Compounds

The prefixes used are mono-, di-, tri-, tetra-, penta-, hexa-, and so forth.

The mono- prefix is usually not used for the first element in the formula.

The "o" and "a" endings of these prefixes are dropped when they are attached to "oxide.”

Also, the ending of the last (most negative) element is changed to -ide.

Four electron pairs around the central atom

The simplest is methane, CH4. Determine its Lewis dot diagram.Tetrahedral: The carbon atom would be at the centre and the hydrogens at the four corners. All the bond angles are 109.5°.Build CH4 & draw it.

An Additional FactorLone pairs are in orbitals that are shorter and rounder than the orbitals that the bonding pairs occupy. Because of this, there is more repulsion between a lone pair and a bonding pair than there is between two bonding pairs.That forces the bonding pairs together slightly - reducing the bond angle from 109.5° to 107°. This is pyramidal.

Water, H2O

Oxygen has four pairs of electrons, two of which are lone pairs. These will again take up a tetrahedral arrangement. This time the bond angle closes slightly more to 104°, because of the repulsion of the two lone pairs.Not tetrahedral, because we only "see" the oxygen and the hydrogens - not the lone pairs. Water is described as bent or V-shaped.

Three electron pairs around the central atom

3 bonds and no lone pairs. The 3 pairs arrange themselves as far apart as possible. They all lie in one plane at 120° to each other. The arrangement is called trigonal planar.

What happens if B is a lot more electronegative than A?

In this case, the electron pair is dragged right over to B's end of the bond. A has basically lost control of its electron, and B has complete control over both electrons. Ions have been formed.

Polar bonds and polar molecules

In a simple molecule like HCl, if the bond is polar, so also is the whole molecule. What about more complicated molecules?In CCl4, each bond is polar.The molecule as a whole, however, isn't polar - in the sense that it doesn't have an end (or a side) which is slightly negative and one which is slightly positive. The whole of the outside of the molecule is somewhat negative, but there is no overall separation of charge from top to bottom, or from left to right.

By contrast, CHCl3 is polar.

The hydrogen at the top of the molecule is less electronegative than carbon and so is slightly positive. This means that the molecule now has a slightly positive "top" and a slightly negative "bottom", and so is overall a polar molecule.A polar molecule will need to be "lop-sided" in some way.