CHEM 205 U1-6

Hydroxide

OH^-

Ammonium

NH4^+

Nitrate

NO3^-

Nitrite

NO2^-

Carbonate

CO3²-

Sulfate

SO4²-

Sulfite

SO3²-

Phosphate

PO4³-

Phosphite

PO3³-

Perchlorate

ClO4^-

Chlorate

ClO3^-

Chlorite

ClO2^-

Hypochlorite

ClO^-

% relative error equation

[ |observed - accepted| / |accepted| ] x 100

precipitation reaction

an insoluble solid forms when substances in solutions react

sig figs x /

same # sig figs as # with least sig figs in expression

sig figs + -

same # decimal places as # with least dec places in expression

(sig figs) if you have an ± before a * or /

use the decimal-place rule for + or − to determine the precision of that result. Then, for × or ÷, use the fewest significant figures among all the numbers (including the answer of + or -). Only round at the very end!

Inaccurate results

could suggest systematic error

Systematic error

consistently shifts measurements in one direction, affects accuracy

Random error

causes unpredictable variation between measurements, affects precision

JJ Thomson

cathode ray experiment → found evidence for the existance of electrons

Robert Millikan

Oil drop experiment → measured electron charge

Ernest Rutherford

Gold foil experiment → proved existence of a small, dense, + charged nucleus in atoms

James Chadwick

Discovered the neutron

Cathode Rays

streams of electrons

Atomic #

# of protons

Mass #

protons + neutrons

Isotopes

Atoms of the same element that have diff # of neutrons

Changing neutrons

changes isotopes

Changing protons

changes element

periods

horizontal rows

groups/families

vertical columns, elements have similar properties

metals (location, etc)

left n center, conductors, malleable, ductile, shiny, usually solid

non metals (location, etc)

mainly right side, can be solid liquid or gas

Metalloids

Right side, dividing staircase between metals n non metals, often semi-conductors

common metalloids

B, Si, Ge, As, Sb, Te

reactivity of metals

lose electrons, reactivity increases as you go down the group

reactivity of non metals

gain/share electrons, reactivity increases up a group

highly reactive metals found

bottom left (alkali)

Highly reactive nonmetals found

Top right (esp halogens) (excluding noble gasses)

Group 2/2A

Alkaline earth metals → highly reactive, +1 ions, occur in compounds

Group 2/2A

Alkali earth metals → reactive, form +2 ions, occur in compounds

Group 17/7A

Halogens → reactive non-metals, form -1 ions

Group 18/8A

Noble gasses → Exist as individual atoms, very unreactive, heavier ones form compounds

Seven diatomic elements

H2, N2, O2, F2, CL2, BR2, I2

the 7 diatomic elements @ room temp

Br2 → liquid, I2 → solid, rest → gasses

H

can form H+ or H-

B and Al from group 3A

B is metalloid, Al forms Al3+

Si is found

in Silicates

N2

nitrogen, relatively unreactive, 78% air

O2

oxygen, supports combustion, 21% air

S is commonly

S8

Allotropes

diff structural forms of the same element

molecular form

An element existing as molecules of 2 or more atoms bonded together (Br2, O2, S8)

element symbol

just the letters (Br)

Valence electrons

determine an element’s bonding n reactivity

valence electrons are

the electrons in outermost shell

Valence electrons in group 1A, 2A, 3A, 4A, etc.

1,2,3,4,etc

exeption to # valence electrons by group

Helium, has 2

Atoms are more stable

when the outermost electron shell is full

the first shell for any atom

Can hold max 2 electrons → full (except hydrogen)

For common atoms the outer shell

Often full with 8 electrons (exept H and He)

Electronegativity

an atoms ability to attract shared electrons towards itself

electronegativity in periodic table

increases up and to the right (noble gasses excluded from this trend)

non metals have higher than metals

electronegativity

Noble gasses don’t need to

loose or gain electrons (already stable)

covalent bonding

atoms share electrons, Non metal + non metal

Ionic bonding

Electrostatic attraction between oppositely charged atoms, (happens with oppositely charged ions AFTER electrons have been transferred), Metal + non metal or compounds containing polyatomic ions such as NH4+

Metallic bonding

Electrons are spread over many metal atoms and can move, metals

sharing in bonding

can be equal or unequal

Molecule

group of covalently bonded atoms

molecular formula

gives the actual # of each type of atom (CH4)

Structural formula

Shows how atoms are connected

Isomers

same molecular formula diff structures

Diff between Atomic and Molecular elements

1) represented by single atomic symbols ( how elements r written in periodic table)

2) represented by formulas containing multiple atoms (O2, N2, S8)

Molecular elements

Molecules containing one element O2, P4, S8

Molecular compound

Molecules containing diff elements H2O, CO2, NH3

Covalent network

Type of solid where a Giant network of atoms are connected by covalent bonds (ex: diamond)

Ionic compound

Bonds of cations and anions

Example of ionic compound

NaCl, MgO

Cations

lose electrons, + charge

Anions

gains electrons, - charge

charge

#protons - #electrons

ion

atom/molcule tht has a net charge cz it lost/gained 1 or more electrons

diff between ionic bonding and ionic compound

Ionic bonding = HOW ions stick together

Ionic compound = WHAT is formed

m³ to L

1 m³ = 1000L

sig fig rules dont apply

conversion factors

density of water

1g/cm³

floats in water

less dense than water (1g/cm³)

sinks in water

more dense than water (1g/cm³)

Group 16

chalcogens

polyatomic ion H2PO4-

Dihydrogen phosphate

polyatomic ion (PO4)³-

Phosphate

polyatomic ion (CO3)²-

Carbonate

polyatomic ion (NO3)^-

Nitrate

chemical names, unless starting a sentence or for Roman numerals

there should be no capital letters.

Fe

Iron

Bromine

Br^-, (2+ or 3+)

Co

Cobalt, (2+ or 3+)