Composition of Matter

Composition of Matter

Matter – anything that has mass and takes up space

Substance – single pure form of matter, not a mixture

Pure – same/uniform throughout even on microscopic scale

The Elements

Elements – fundamental substances from which all matter is built – matter is combination of elements.

Chemical elements are the basic building blocks of matter and in various combinations make up all the matter on earth

Atom – smallest particle of an element that exists

Element – is a substance composed of only one kind of atoms.

Names of the elements

Each element (~116) has a name and unique chemical symbol

Gold, Au (Old English word meaning ‘yellow’ Au, aurum)

Characteristic property – Chlorine – yellow-greenish

Persons, geographical place etc. English, Greek, Latin, German…

Chemical symbol – 1 (capital) or 2 (1st capital, second small) letters

a. First one or two letters of element name: H, C, N, O, He (not HE), Ne

b. First and later letter of name: Zinc –Zn, Magnesium – Mg

The Nuclear Atom

Current summarized model:

i. Atoms are made up of subatomic particles called electrons, protons and neutrons.

ii. The protons and neutrons form a compact, central body called the nucleus of the atom.

iii. The electrons are distributed in space like a cloud around the nucleus.

Particle

Symbol

Charge*

Mass, g

Electron

e

-1 9.109 x 10-28

Proton

p

+1

1.673 x 10-24

Neutron

n

0

1.675 x 10-24

(* 1.602 x 10-19 Coulomb)

Positive charge of protons (in nucleus) cancels negative charge of electrons (outside nucleus) – so atoms are electrically neutral.

Electrons – J. J. Thomson (cathode rays…)

Rutherford, Geiger, Marsden, ‘Solar system’ atomic model

Number of protons in an atom is the atomic number, Z of that element.

Atomic number is the characteristic that distinguishes one element from another

H, Z = 1; He, Z = 2; Au, Z = 79 etc

Neutrons and protons are jointly called nucleons.

Total number of nucleons (protons + neutrons) in nucleus is the mass number, A of the atom.

Example:

Determine number of electrons in atom with A = 238 and 146 neutrons

A = p + n = 238

N = 146

P = A – n = 238 – 146 = 92

But p = e in an atom, so the number of electrons = 92

Neutrons, with no charge, do not affect number of electrons (or Z) of an atom, but do affect the mass and mass number.

Atoms with same atomic number (Z) but different mass number (A) are isotopes of that element. Isotopes of same element have same atomic number (same p, same e) but different number of neutrons in the nucleus.

Naming isotopes: element-mass number

Eg. Neon-20; Uranium-235

Symbol:

EMBED ChemDraw.Document.6.0

X

Z

A

Ex: Number of neutrons in nucleus of

Sr

90

38

# Neutrons = A – Z 90 -38 = 52

Ex: Write nuclear symbols for isotopes of oxygen having 8, 9 and 10 neutrons respectively

O,

O,O

16

1718

8

8

8

Isotopes of same element have same chemical, and similar physical properties (chemical properties determined by Z), except for hydrogen 1H; Deuterium (D) 2H; tritium (T) 3H; have such large differences in mass numbers that have noticeable differences in some physical properties

Ex: Volume of 100 g of water and 100 g of heavy water

The Periodic Table

Listing of all elements according to increasing Z and arranged so that elements with similar chemical/physical properties occur in vertical columns (family or groups); location gives general idea about chemical/physical properties

The arrangements of elements that shows their family relationships is called the periodic table

1, 2, 13 -18 Main group representative elements

Group 1 (IA) most active metals, M+

Group 2 (IIA) alkaline earth metals

Group 17 (VIIA) halogens – most active non-metals

Group 18 (VIIIA, 0) inert/noble gases

Group 3 – 12 transition metals/elements

Lanthanides

Actinides

Metals

Non-metals

Metalloids

Diagonal: metalloids (sometimes); (B), Si, Ge, As, Sb, Te Po

Metals – lower, left ; conduct electricity, metallic luster, malleable (hammered into sheets) and ductile (drawn into wires)

Non-metals – upper, right ; Electrical insulator, brittle solids, gases

Metalloids – diagonal – semiconductors, physical properties of metals, but chemical properties of non-metals

Compounds

Most substances are combination of elements rather than pure elements

Compounds – a substance that consists of two or more elements in definite ratios by mass – law of constant composition: fixed ratio by mass; Water, H2O always 88.8% oxygen, 11.2% hydrogen by mass

Organic compounds – compounds containing C and usually H too. Examples: methane, propane, sugars etc. (usually molecular)

Inorganic compounds – all other compounds, molecular or ionic (includes CO2, CO, CO32- etc.)

The elements in a compound are not just mixed together.

Atoms in compounds are joined or bonded to each other – covalent bonds in molecules or ions in ionic solids.

Molecules – definite, electrically neutral group of covalently bonded atoms – H2O, NH3, CH4 etc. Molecular compounds have molecules; most organic compounds are molecular; room-temperature gases, liquids are molecular.

Ions – positively (+ve) or negatively (-ve) charged atoms or groups of atoms; Na+, Cl, NO3, SO42 etc.

Ionic solids consist of ions. Ionic compounds tend to be high-melting solids

Molecular Compounds

Chemical formula – chemical symbols for atoms in compound

Molecular formula – Chemical formula for molecules showing numbers and types of atoms in the molecule:

Water, H2O – two H, one O atom per molecule

Sugar, Sucrose C12H22O11, 12 C, 22H and 11O atoms

Molecules have geometry and specific arrangement of atoms

Structural formula

Represent the atoms by symbol and use lines to show which atoms are joined together.

CH

H

H

COH

H

H

Ethanol

Shows atoms, bonds but not geometry

Ball-and-stick model – shows geometry: bond angles and lengths

Space-filling model – atoms represented by spheres that fit into one another. Computer generated from MO calculations.

Ionic compounds

Ionic compounds consist of positive and negative ions held together by electrostatic forces in a geometric array (crystal lattice).

cation – positive ion (Na+)

anion – negative ion (Cl)

Monoatomic ions – single atoms that have gained or lost electrons

In an atom, number of electrons = number of protons, so net charge = 0

Loss of electron leaves excess positive charge ( cation

Na (atom) ( e + Na+ (cation)

Loss of several electrons:

Ca (atom) ( 2 e + Ca2+ (cation)

Al (atom ( 3 e+ Al3+

Gain of electron(s) causes excess negative charge ( anion

Cl (atom) + e ( Cl (anion)

O (atom) + 2 e ( O2 (anion)

N (atom) + 3 e ( N3 (anion)

Metals typically form cations, nonmetals form anions; common ions often related to group numbers:

Cations: 1, 2: M+, M2+; 13 (13-10) = M3+, Al3+

Anions : group 17, halogens, 17 – 18 = -1

Group 16, oxygen group, 16 – 18 = -2

Group 15, nitrogen group 15 – 18 = -3

Transition metals: 3 – 12, multiple positive charges

Heavy group 13 -15 elements: two cations, Pb2+ and Pb4+ etc

Example:

Probable ions for radium (Ra), astatine (At), tellurium (Te), zinc (Zn)

Ra, group 2, Ra2+

At, group 17, non-metal, 17 – 18 = -1, At

Te, group 16, metalloid, 16 – 18 = -2, Te2

Zn, group 12, transition metal (memorize as) Zn2+

(All compounds are electrically neutral overall)

Polyatomic ions – several atoms covalently bonded together as a unit with net positive or negative charge: NH4+, SO42

NH

H

H

H

+

Protons 7 + 4(1) = 11;electrons 11 – 1;net charge = +1

Oxoanions – polyatomic anion with oxygen atom(s) around another central atom, CO32, carbonate; NO3, nitrate; PO43 phosphate

(Refer text book and handout)

Ionic compounds have many cations electrostatically bonded to many anions, so do NOT contain discrete molecules. Ratio of number of cations to number of anions serve as formula for the compound.

Chemical formula – smallest, whole number ratio of cations and anions.

Examples: One Na+ per one Cl in NaCl

Two Na+ per one CO32 in Na2CO3

Ionic compounds must be electrically neutral, so use ion charges to ‘balance/write’ formulas:

Cerium(III)sulfate

Ce3+ and SO42- ; 2(+3) + 3(-2) = 0, Ce2(SO4)3

Cerium(IV)sulfate, Ce(SO4)2

(Do not write ion charges in actual formula)

Formula unit – smallest ‘unit’ of ionic compound – same number of atoms as formula

Formula units: NaCl, (NH4)2SO4

Example: Calculate ratio of atoms in mica, KMg3Si4AlO10(OH)2.

Al:H:K:Mg:O:Si = 1:2:1:3:12:3

(Mixtures/Separation of mixtures – Please read pp 25 – 30 (4th edn.) 21 – 26 (3rd edn.).)

Nomenclature (naming) of compounds

Common names vs systematic names

Cations – Monoatomic cations names as element plus ‘ion’

Na+, ‘sodium ion’, stock number (roman numeral identifying the charge) is included for elements with more than one (common) ionic charge: Co2+, Co3+, Co+, Cobalt(II), Cobalt(III), Cobalt(I); needed for transition metal, heavier group 13 -15 metals.

Anions

1. Monoatomic anion named by adding ‘ide’ to stem of element name plus ‘ion’

F fluorine ( fluoride ion

Cl chlorine ( chloride ion

X halogen ( halide ion

O2, oxygen ( oxide ion; N3 nitrogen ( nitride ion

2. Oxoanions – anions with oxygen around another central atom

a. Named by adding ‘ate’ to stem of name of central atom

CO32 - carbonate; SiO44 silicate

b. If two oxoanions of same central atom -

- use ‘ate’ for oxoanion with more O atoms

- use ‘ite’ for oxoanion with fewer O atoms

NO3, NO2 nitrate, nitrite

SO42, SO32 sulfate and sulfite

c. If four oxoanions of same central atom (halogens)-

- use prefix ‘hypo’ and ‘ite’ ending for least O atoms

- use ‘ite’ and ‘ate’ ending for 2nd and 3rd least O atoms, respectively

- use prefix ‘per’ and ‘ate’ ending for most O atoms

ClO - hypochlorite

ClO2 - chlorite

ClO3 - chlorate

ClO4 - perchlorate

Oxoacids are molecular compounds, ‘parent’/sources of oxoanions.

Obtain acid by adding H+ to oxoanions.

Names:

I. The ‘-ate oxoanions gives ‘-ic’ acid

II. The ‘-ite’ oxoanions gives ‘-ous’ acid

HClO2

chlorous acid

HClO3

chloric acid

HClO

hypochlorous acid

HClO4

perchloric acid

H2SO4

sulfuric acid

HNO3

nitric acid

H3PO4

phosphoric acid

Binary acids – ‘hydro – ic’ acid –IN WATER solution

HCl(aq)hydrochloric acid; HCl(g) hydrogen chloride

H2S(aq)hydrosulfuric acid; H2Shydrogen suflde

3. Some anions include H in formual and name:

HCO3

hydrogen carbonate or bicarbonate

HPO42hydrogen phosphate

H2PO4dihydrogen phosphate

Naming ionic compounds

Name cation(s) first, anion(s) next as separate words, omit ‘ions’

KCl

potassium Chloride

NaH2PO4sodium dihydrogen phosphate

CuCl2

copper(II) chloride

ZnBr2

zinc bromide

Hydrates – ionic salts containing H2O molecules as part of the crystal/formula

CuSO4.5H2O, CoCl3.6H2O

Name – Greek prefix (for number) + ‘hydrate’

1mono-

5penta-

9nona-

2di-

6hexa-

10 deca-

3tri-

7hepta-

11undeca-

4tetra-

8octa-

12dodeca-

CuSO4.5H2O

Copper(II)sulfate pentahydrate

CoCl3.6H2O

cobalt(III) chloride hexahydrate

Naming Molecular Compounds

-use Greek prefix for number of atoms, unless one;

-write less electronegative atom first, more electronegative element (group 16, 17) second, with name ending in ‘ide’

N2O5dinitrogen pentoxide

SF6sulfur hexafluoride

N2Odinitrogen oxide

Exception COcarbon monoxide

P2O5 (NOT phosphorus (V) oxide), the molecule is P4O10 named as tetraphosphorus decoxde.

Common names: NH3 ammonia, N2H4 hydrazine; PH3, phosphine; NH2OH hydroxylamine

Example:

N2O4

dinitrogen tetroxide

P4S5

tetraphosphorus pentasulfide

Formulas from Name –

a. Identify compound as ionic or molecular – metals tend to form ionic compounds, two nonmetals are usually molecular

b. Ionic: Symbol of metal (cation) first, nonmetal or polyatomic anion second; balance charges with subscripts (number of ions). Recognize charge by group position or stock number.

c. Molecular: Write element symbols in order names, using Greek prefixes for number of atoms

Example:

Magnesium nitride, ionic: Mg, 2, Mg2+

N, 15, 15-18 = -3 N3-; 3(+2) + 2(-3) = 0 Mg3N2

Diboron trisulfide, molecular: B2S3

Potassium iron(III) sulfate dodecahydrate (alum), ionic:

K+, Fe3+, SO42-, 12 H2O, +1+3+2(-2) = 0

KFe(SO4)2.12H2O

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