Mast and Hurley Chap 2.ppt

Atoms, Elements and Ions

Boyle - Early Measurements

•Worked with gases (Boyle’s Law)

Elements

• May be reduced to a single atom

• 88 occur naturally

• Others created in particle accelerators and are short-lived

• Periodic Table to 118 but not all elements

Symbols for Elements

• Symbolic representations standard - consists of a single capital letter OR a capital letter and a

lowercase letter.

• The letters are usually the first, or first and second, letter in the current (or Latin) name.

Law of Constant Composition

• Compounds consist of combinations of elements; • Different compounds always have the same

proportion (by mass) of the elements.

Dalton’s Atomic Theory

Developed theory on the Law of Constant Composition. 1. Elements are made of tiny particles called atoms 2. All atoms of a given element are identical

3. Atoms of a given element are different from other elements 4. Atoms of different elements can combine --> compounds 5. A given compound always has same type/number of atoms 6. Atoms are indivisible in chemical processes - they are not

Atomic Structure

• Thomson (1890’s) - atom contains negative particles (electrons) • 1910 - Lord Kelvin’s Plum Pudding model (electron (-) charges

in a (+) cloud)

• Rutherford, showed the plum pudding model to be bogus by bombarding metal foil with alpha particles; some particles were deflected (alpha particles are the result of radioactive decay…which we may talk about at the end of the year)

Rutherford’s Atomic Model

• He concluded the center of an atom must be dense and small - Nucleus

• The electrons are (-) and move through a space surrounding the nucleus

• He concluded the nucleus includes protons; a proton

is (+) to balance the negative electrons.

Modern Atomic Model

• Consists of three subatomic particles

– Protons: mass = 1.0073 amu, charge = +1 – Neutron: mass = 1.0087 amu, charge = 0 – Electron: mass = 5.486 X 10-4, charge = -1

An amu (atomic mass unit) = 1.66054 X

10

grams.

Modern Atomic Model

• Atoms are extremely small

– Most have diameters between 1 X 10-10 m and 5 X 10-10 m

– Diameter often expressed in angstroms (Å) 1 Å = 10-10 m

• Charge often expressed in coulombs

Isotopes

•

All atoms of a particular element

have the same number of protons.

(The Atomic Number)

• Atoms of the same element with

differing numbers of neutrons are

isotopes

.

• An atom of a specific isotope is called a

Isotopes

• Example

11

C: 6 protons, 6 electrons, 5 neutrons

C: 6 protons, 6 electrons, 6 neutrons

C: 6 protons, 6 electrons, 7 neutrons

C: 6 protons, 6 electrons, 8 neutrons

Atomic Weights

Atomic mass is also known as its atomic weight

• Atomic particles have mass

• Compounds have mass

• Mass in reactions are maintained

•

Atomic mass represents the relative

Symbols for Elements

2

He

Calculating Average Atomic Mass

Naturally occurring chlorine is 75.78% 35Cl (atomic mass 34.969 amu) and 24.22% 37Cl (atomic mass of 36.966 amu). Calculate the average atomic mass (atomic weight) of chlorine.

Average atomic mass =

(0.7578)(34.969 amu) + (0.2422)(36.966 amu) = 26.50 amu + 8.953 amu

Chemical Bonds

• Are formed between atoms

• Hold Atoms of compounds together

• Are formed to achieve stabile e

-configurations

• Are of two types

– Ionic

Simple Formula’s

• Are for a single molecule

• Show number and type of atoms

Examples

CaCl

(1 Calcium, 2 Chlorine atoms)

Types of Ions

Anion

- Have gained electrons

S

, Cl

Cation

- Have lost electrons

Ionic Compounds

• Form because cation and anions attract

each other

• Are held together by IONIC BONDS

(electrostatic attraction)

• Must be balanced

Examples

Periodic Table of the elements

•

http://www.ptable.com/

•

http://www.webelements.com/

• Note

– Atomic Radius

The Periodic Table

• Is arranged by atomic number (Mendeleev arranged this way for the first time)

• Is organized into groups (families) of elements

• Reflects the atomic organization of electrons

Metals

• Group 1 Metals: Alkali Metals (not H)

– extremely reactive, soft and lustrus, never found as pure element

• Group 2 Metals: Alkaline Earth Metals

– Very reactive, soft, not found as pure elements

• Transition (Groups 3-12) and other Metals (found in Groups 13-16)

Atoms as Ions

• Atoms that lose, or gain, extra electron(s) become charged

• Electrons occupy orbitals surrounding the nucleus that contain specific numbers of electrons

• Atoms lose, or gain, electrons to become more stable energetically

– An empty, half-full, or full orbital is more stable than other proportions

Non-metals

• Poor conductors, not lustrus, brittle or powders in solid form

• Noble Gases (Group 18)

– are non-reactive and gaseous at room temp

• Halogens (Group 17)

– highly reactive - gaseous or liquid (Br) or solid (I, At) at room temp; diatomic (e.g. Cl₂)

• Other Non-Metals

Metals

• Are good conductors, mallable and ductile

• Are generally lustrus (shiny)

• Donate Electrons in reactions

• Form salts with non-metals

Metalloids

• Includes B, Si, Ge, As, Sb, Te

• Conductivity between that of metals and

non-metals

• Important in computers

Atomic Radii

• Size of the atom…

• Increases as go down the table

Electron Affinity

• Energy change that occurs when an electron is added to a gaseous atom or ion.

• If the addition is exothermic, the sign is (-) • X _(g) + e- --> X- _(g)

• The more negative, the more energy is released.

• Generally, as go across table, electron affinity increases and as go down, it decreases

Electronegativity

• Ability of an atom in a molecule to attract shared electrons to itself.

• Increases across table, decreases down the table.

• Fluorine the most electronegative

• Greater electronegativity = greater attraction for electrons

• Useful when considering bonding types between atoms and solubilities…

Covalent Bonds

• Represent bonds in which two atoms share electrons

• Formed between nonmetals (molecular compounds)

– Example: CO₂, C₂H₄

• Formed between diatomic atoms

Ionic Bonds

• Represent reactions in which electrons

have been transferred

• Are based on electrostatic attraction of

ions

– Anion: negative ion (O-2) – Cation: positive ion (Na+1)

Ionic Compounds

• Are always electrically neutral

• Writing formulas

– balance the charge

– If the charge isn’t equal on atoms, the

charge on one atom (without the sign) will become the subscript on the other

Example:Mg+2 and N-3 Mg₃N₂

Naming Positive Cations

Cations formed from metal atoms have

the same name as the metal

Naming Cations

If a metal ion can form cations of differing charges, the positive charge is given a Roman numeral in parentheses following the name of the metal

Examples

– Fe2+ is the iron (II) ion – Fe3+ is the iron (III) ion – Cu+ is the copper (I) ion – Cu2+ is the copper (II) ion

Naming Cations

Older method of naming ions of same metal with different charges

– Fe2+ is the ferrous ion – Fe3+ is the ferric ion – Cu+ is the cuprous ion – Cu2+ is the cupric ion

Uses the old latin name and -ous suffix to lower charge, -ic to higher charged ions

Naming Cations

Cations formed from nonmetal atoms

have names that end in -ium

Examples

Common Cations

Naming Anions

Monatomic Anions have names formed by replacing the ending of the name of the element with -ide

Examples

H- hydride ion O2- oxide ion S2- sulfide ion Cl- chloride ion

Same for some simple polyatomic ions

Naming Anions

Polyatomic anions containing oxygen (or oxyanions) have names in -ate or -ite

-ate indicates most common anion

-ite indicates ion with same charge but 1 fewer oxygen atom

Examples

Naming Anions

Prefixes are used when the series of oxyanions of an element extends to four members, as with the halogens

Example

ClO₄- perchlorate ion (per indicates 1 oxygen atom) ClO₃- chlorate ion

ClO₂- chlorite ion (one less atom than chlorate)

Naming Anions

Anions derived by adding H+ to an oxyanion are

named by adding as a prefix the word hydrogen or dihydrogen (as appropriate)

Examples

CO₃2- carbonate ion PO₄3- phosphate ion With hydrogen(s)…

HCO₃- Hydrogen carbonate ion

Common Anions

Naming Ionic Compounds

Names consist of the cation name followed by the anion name

Examples

CaCl₂ calcium chloride Al(NO₃)₃ aluminum nitrate

Cu(ClO₄)₂ copper (II) perchlorate (or cupric perchlorate)

Acids and Bases

• Contain ionic bonds

Bronsted-Lowry Model

• Acids

– Are proton (H+) donators

• Bases

Acids/Bases

Common and easy to understand Acids include

HCl --> H+ + Cl- (dissociation equation in solution)

H₂SO₄ --> H+ + HSO₄ (dissociation equation in solution)

H₂SO₄ --> 2H+ + SO₄ (dissociation equation in solution)

Common and easy to understand bases include

NaOH --> Na+ + OH- (dissociation equation in solution)

Mg(OH)₂ --> Mg2+ + 2OH- (dissociation equation in solution)

Naming Acids

• If the anion does NOT contain Oxygen

– Prefix hydro added to the front

– suffix -ic added to root name of element

Examples

Naming Acids

• When the anion CONTAINS oxygen

– Acid name formed from root name of central element of the anion or the anion name

– Suffix of -ic or -ous added

Examples

H₂SO₄ Sulfuric Acid (anion is sulfate)

H₂SO₃ Sulfurous Acid (anion is sulfite)

H₂CO₃ Carbonic Acid (anion is carbonate)

Molecular Compounds

Compounds with only Nonmetals

• Covalent bonds are between atoms

– Sharing of electrons holds atoms together – There are no electrostatic charges

– Compound may be polar if one atom gets “more of the share”

Naming Molecular Binary Compounds

The rules…

• First element in formula named first

• Second element named as if an anion

• Prefixes used to denote number of

atoms present

Prefixes

Prefix Number indicated

Examples…

• N

O dinitrogen monoxide

• NO

nitrogen monoxide

• N

O

dinitrogen pentoxide

(Note, the vowel of the prefix is dropped