Review Unit

Investigation Report Outline

Purpose

Most often to create, test or use a chemistry concept.

Problem

A specific question to be answered in terms of the manipulated and responding variables.

Hypothesis and/or Prediction

A predicted answer to the problem, including reasoning.

(pgs. 790 – 794)

Investigation Report OutlineMaterials

Complete list of all equipment and chemicals.

Procedure

Detailed set of numbered steps to answer the Problem.

Evidence

All quantitative and qualitative observations relevant to answering the Problem. Use tables whenever possible.

Investigation Report Outline

Analysis

Manipulations, interpretations and calculations based on the evidence. It concludes with a statement that answers the Problem.

Evaluation

1. Evaluation of the Investigation

2. Evaluation of the Prediction

3. Revisiting the Purpose

pgs. 4 & 5 Are You Ready #’s 1, 2, 4, 5, 9, 10, 11

Pg. 6 Starting Points #’s 1 – 5

Read pgs. 790 - 794

Read pgs. 6 – 11 pg. 11 Section 1.1 Questions #’s 1 – 8

Semi-metals (or metalloids) are all of the elements that border the staircase line except aluminium.

Read pgs. 12 – 13

pg. 13 Section 1.2 Questions #’s 1 – 5

Read pgs. 14 – 16

pg. 16 Section 1.3 Questions #’s 1 – 8

John Dalton(1766 - 1844)

The “billiard ball” model

Dalton’s Atomic Theory

Matter is composed of indestructible, indivisible atoms which are identical for one element, but different from other elements.

J.J. Thomson(1856 – 1940)

The “raisin bun” model

An atom consists of one large positive charge and many small negative charges embedded in it.

Thomson’s Atomic Model

"Could anything at first sight seem more impractical than a body which is so small that its mass is an insignificant fraction of the mass of an atom of hydrogen? --which itself is so small that a crowd of these atoms equal in number to the population of the whole world would be too small to have been detected by any means then known to science."

- an excerpt from a speech he made in 1934J.J. Thomson(1856 – 1940)

Ernest Rutherford(1871 –

1937)

The “planetary” or “nuclear” model

Every atom has a tiny, extremely dense positive core which he called the nucleus.

The negative electrons orbit the nucleus like planets around the sun.

Rutherford’s Atomic Theory

Rutherford was surprised when all the particles did not go straight through the gold foil. He realized that each atom must have a dense core of positive charge.

The mass number of an atom is equal to the number of particles in the nucleus (protons and neutrons).

The atomic number of an atom is equal to the number of protons in the nucleus.

Three naturally occurring isotopes of carbon:

126 C 13

6 C 146 C

The bottom number is sometimes not written because you can determine the atomic number from the symbol.

Any sample of an element found in nature is a mixture of different isotopes. Each isotope will occur in different proportions, usually given as a percentage.

For example:Each of these isotopes contains 50 protons.

Neils Bohr(1885 –

1962)

Electrons are arranged around the nucleus in very specific orbits or energy levels.

Bohr’s Atomic Theory

The period number (horizontal row) that an element is in is the

same as the number of energy levels the atom has.

Level Max. # of Electrons

1 2

2 8

3 8

4 18

All the elements in a group have the same electron configuration in their outermost shells. Electrons in the outer shell that is not full are called valence electrons.

When an electron absorbs energy, it jumps to a higher energy level.

When an electron transitions back to a lower level, energy is given off.

The light given off by the different possible transitions is called a line spectrum.

A different colour is emitted for each different transition.

Every different element has its own characteristic emission spectrum.

Bohr used the light emitted from excited hydrogen atoms to form his ideas about the structure of the atom.

The hydrogen emission spectrum

An ion is a an atom (or a group of atoms) that has a positive or negative electric charge.

Recall that an atom of any element is neutral, so the number of protons equals the number of electrons.

The formation of an ion is called ionization, and is the result of an atom either gaining or losing electrons.

The number of protons only changes in nuclear reactions,

never in the formation of ions.

Cations are positively charged ions.

Ant-ion?

Anions are negatively charged ions.

They are formed when a metal atom loses valence electrons (electrons in the outermost energy level).

They are formed when a non-metal atom accepts electrons into its outer energy level.

Read pgs. 18 – 25

pg. 26 Section 1.4 Questions #’s 1 – 15

• ionic compounds: metal + non-metal

• molecular compounds: non-metal + non-metal

Empirical Classification of Ionic and Molecular Compounds

Recall that empirical knowledge is observable.

Names and Formulas of Ionic Compounds

Binary Compounds

aluminum fluoride

silver sulfide

potassium iodide

3 sAlF

2 sAg S

sKI

sMgO

3 2 sCa P

2 sBaBr

You must ALWAYS include the state of matter when writing an empirical formula.

magnesium oxide

calcium phosphide

barium bromide

Names and Formulas of Ionic Compounds

Multi-Valent Metals

iron (III) chloride

lead (IV) oxide

nickel (III) sulfide

3 sFeCl

2 sPbO

2 3 sNi S

2 sFeCl

sCrN

2 sCu O

If the ion of a multivalent metal is not specified, you can assume the charge on the ion is the most common one.

iron (II) chloride

chromium (III) nitride

copper (I) oxide

The most common ion charge is the one that is listed first.

Names and Formulas of Ionic Compounds

Compounds with Polyatomic Ions

barium hydroxide

iron (III) carbonate

copper (I) permanganate

gold (III) nitrate

ammonium phosphate

potassium dichromate

3 3 (s)Au NO

4 4 (s)3NH PO

2 2 7 (s)K Cr O

2 (s)

Ba OH

2 3 3 (s)Fe CO

4 (s)CuMnO

The polyatomic ion must be in brackets only when there is more than one.

Ionic Hydrates

An ionic hydrate is an ionic compound that has loosely-bonded water molecules.

4 2 (s)CuSO 5H Ocopper (II) sulfate pentahydrate

copper (II) sulfate-5-water

copper (II) sulfate-water (1/5)

Read pgs. 27 –31

pg. 32 Section 1.5 Questions #’s 1 – 17

water ( H2O(l) )

ammonia ( NH3 (g) ) methane ( CH4 (g) )

A molecule is a group of nonmetal atoms held together by covalent bonds.

A molecular formula indicates the number of atoms of each type.

S8 (s)

P4 (s)

H2 (g)

Br2 (l)

Molecular Elements

“I Bring Clay For Our New House.”

2 (g) 2 (l) 2 (g) 2 (g) 2 (g) 2 (g) 2 (g)I Br Cl F O N H

4 (s) 8 (s)P S

“And four Paving stones for eight Steps.”

“And three Octopuses for pets!”

3 gO

In a covalent bond, electrons are shared between nonmetal atoms in pairs.

two chlorine atoms one chlorine molecule

a pair of shared electrons

2 gCl

water (H2O)

an oxygen atom and two hydrogen atoms

a water molecule

two pairs of shared electrons

In a covalent bond, electrons are shared between nonmetal atoms in pairs.

Name Formula

carbon dioxide

dinitrogen monoxide

phosphorus trichloride

oxygen difluoride

dinitrogen tetrasulfide

sulfur trioxide

2 (g)CO

2 (g)N O

3 (g)PCl

2 (g)OF

2 4 (g)N S

3 (g)SO

PrefixNumber

of Atoms

mono 1

di 2

tri 3

tetra 4

penta 5

hexa 6

septa 7

octa 8

nona 9

deca 10

Naming Binary Molecular Compounds

When there is one of the first element only, then the prefix mono is dropped.

The common names of these compounds are to be memorized.

Naming and Writing Formulas for Acids and Bases

Formula IUPAC Name Classical Name

aqueous hydrogen bromide

aqueous hydrogen chromate

aqueous hydrogen nitrite

aqueous hydrogen sulfate

aqueous hydrogen iodide

aqueous hydrogen sulfite

hydrobromic acid

chromic acid

nitrous acid

sulfuric acid

hydroiodic acid

sulfurous acid

(aq)HBr

2 4 (aq)H CrO

2 (aq)HNO

2 4 (aq)H SO

(aq)HI

2 3 (aq)H SO

To name acids, we use the three rules on the data booklet. Note that all acids are aqueous solutions.

Read pgs. 33 –36

pgs. 36 –37 Section 1.6 Questions #’s 1 – 10