Download Matter: A) Homogeneous Matter • Uniform and in 1 phase • Even

Matter:
A) Homogeneous Matter
 Uniform and in 1 phase
 Even distribution of material throughout the sample
1)
Homogeneous Mixture (Solution)
Salt + Water
 Properties of materials vary
 You can separate the components physically (evaporation, distillation)
2)
Pure Substance
a) Elements
 Simplest form of matter
 Made up of one kind of atom
 Can’t be decomposed chemically
I) Metals
 Luster
 Malleability (Strong Bonds)
 Ductile (Ability to draw into wire)
 Good conductors of heat/electricity
 Hard
 High melting points
II)
Non-Metals
 Dull
 Poor conductors
 Brittle
 Low melting point
b) Compounds: 2 or more different elements chemically combined and chemically decompose
B) Heterogeneous (Mixtures)
Sand + Water
 Not uniform and more than 1 phase
 Can separate physically (filtration/sifting)
Properties:
Physical: Observed without producing a new substance
Extensive:
Depends on amount of sample
(mass, volume, weight)
Intensive:
Doesn’t depend on amount of sample
(Density, boiling/freezing point, pH, odor, taste, color)
Chemical: Reacts or doesn’t react with another substance
Reaction leads to chemical change and a new product with its own properties
(Rust, burning, cooking, respiration, digestion)
Kilo
K
1,000 (103)
Deci
D
.1 (10-1)
Centi
C
.01 (10-2)
1 dm3 = 1L
Milli
M
.001 (10-3)
Micro
Ч
.000 001 (10-6)
1 cm3 = 1 mL
Nano
N
.000 000 001 (10-9)
Angstrom (not metric, but
Å
.000 000 000 1 (10-10)
measures atomic dimensions)
Pico
P
.000 000 000 001 (10-12)
Precision: How clustered together the data points are
Accuracy: How close the data points are to the accepted value
Force – Newtons
Pressure – Pascals
Energy – Joules
Mass – Kilogram
Temperature – Kelvin
Amount of substance – Mole
Chemical Reaction: A process by which one or more substance is changed into one or more different substance with a different set
of properties and is represented by the following equation:
Reactant(s)
Product(s)
Conservation of Charge: Na+1 Cl-1
Conservation of Mass: Coefficients are used to balance the number of atoms on each side of the equation
The Seven Diatomics:
H2
F2
N2
Cl2
O2
Br2
I2
*H2S is gaseous
Types of Reactions:
Synthesis (Direct Combination): Element/compound + element/compound  One compound
2 Mg + O2  2 MgO
Decomposition (Analysis): One compound  Element/compound + element/compound
CO3 breaks down into CO2
CuCO3  CuO + CO2
Single Replacement: Element + compound  Compound + element
Cu +Fe(NO3)2  Cu(NO3)2 + Fe
Double Replacement (Exchange of ions): Compound + compound  Compound + compound
HCl + NaOH  NaCl + H2O
Combustion: Hydrocarbon + Oxygen  Carbon Dioxide + Water
CH4 + O2  CO2 + H2O
Water is always gaseous during a combustion reaction
Exception: 2 CH3OH + O2  4 H2 + 2 CO2
Alkali Metal + Water  Hydrogen gas + Base
2 Na + 2 H2O  H2 + 2 NaOH
Solubility rules determine phases
Cation: Positive charged ion (gets smaller, are metals, and loses electrons)
Anion: Negative charged ion (gets larger, are nonmetals and gains electrons)
History of the Atom:
Democritus: World is made of empty space and tiny particles “atomos”
Aristotle: Matter is continuous (not made of small particles)
False
Boyle: Worked with gases (V vs. P) and was the father of modern chemistry.
Lavoisier: The law of conservation of mass (matter can’t be created or destroyed, it can only be rearranged)
Mendeleev: Founded the periodic table
Proust: Law of definite proportions (substances contain the same ratio by mass) and law of constant composition stating that matter
could not be made or destroyed in a chemical reaction. (No matter how you make H2O, it always has the same ratio by mass)
Ex: 2 H2 + O2  2 H2O
2 H2O2  2 H2O + O2
*Both have 1 part mass H: 8 part mass O
Dalton: Developed atomic theory, law of multiple proportions, law of partial pressures, and believed the atom was still a solid
indivisible sphere.
Law of Multiple Proportions: When elements combine to form different compounds, they are in a small whole number ratio
SnO
Sn = 118g
O = 16g
SnO2
Sn = 118g
O = 32g
The oxygens are in a 1:2 ratio with each other
Atomic Theory:
1)
2)
3)
4)
Everything consists of small, indivisible particles that can't be destroyed or created, known as atoms
Atoms of the same element are identical (he didn't know about isotopes) and atoms of different elements are different
Compounds form when different elements chemically combine
Chemical reactions involve the reorganizing of atoms in small whole number ratios
Gay-Lussac: Under constant conditions, volume of gases and products are in small whole number ratios


2 H2 + O2  2 H2O
2 volumes + 1 volume
3 volumes
Avogadro: Equal volumes of gas under same temperature and pressure contain same number of particles/molecules etc.
Parts of the Atom:
Thomson: Cathode ray tube (Crookes Tube) determined the electron’s charge to mass ratio and its deflection. The electrons move
from the negative to cathode to the anode. Created the "Plumb Pudding" model. Small electrons were embedded a the positive
sphere
Planetary Model: Atomic model with electrons revolving around the nucleus. Was created by Rutherford and perfected by Bohr
Millikan: Oil drop test found mass of electron to be 1/1830th of hydrogen (proton). He used an atomizer to make the X rays interact
with the oil droplets where it became charged with electrons
Rutherford: Gold foil experiment with alpha particles (positive) bombarding a piece of gold. Most of them went through the foil, but
some were deflected. Shows that the atom is mostly empty space and has a small, positive, dense center (nucleus). It disproved
Thomson. Also discovered alpha, beta, and gamma radiation.
Orbital Model: Electrons don't follow specific paths, but move freely around the nucleus in areas called orbitals
Chadwick: Alpha Particle  Beryllium  Neutron (High energy, no charge, mass equivalent to proton)
Wave-Mechanical Model: Electrons are located in an area of highest probability based off Heisenberg's Uncertainty Principle
(position and momentum can't be known)
Isotopes: Same atomic number but different mass number due to a difference in neutrons
Nuclide: Nucleus of a specific isotope

- Proton
– Deuterium
- Tritium
Atomic Number (Z): Number of protons in an atom (found on the bottom left of the symbol) (found on top of modified table)
Mass Number (A): Number of protons + number of neutrons (found on top left of the symbol)
Atomic Weight/Average Atomic Mass: An average of all naturally occurring isotopes (found on bottom of modified table)
Mosely: X-rays wavelength depends on proton
1/12th the mass of C-12 = 1 amu or u
Becquerel: Matter containing uranium exposed photographic film
Pierre and Marie Curie: Uranium and radium
Periodic Table
Group/Column/Family: Vertical section with elements of similar properties
Periods/Row: Horizontal section based on increasing atomic number
Group 1: Alkali Metals
Group 2: Alkaline Earth Metals
Group 6: Chalcogens
Group 7: Halogens
Group 8: Noble gases/Stable octet/Inert
Empirical Formula: Simplest whole number ratio of atoms (always used in ionics)
Molecular Formula: Tells you how many atoms are actually present
Structural Formula: Shows how the elements are bonded together
Ionic: Contains a metal and nonmetal. Large electronegativity difference. Transfers electrons. Forms salts (which are usually solids)
Covalent: Contains two nonmetals with little or no electronegativity difference. Shares electrons
Azide: N3-1
Metals: Positive ions immersed in a sea of mobile electrons
Review of Naming
Ionic:
 Always written with empirical formula (lowest whole number ratio)
 Name the metal and then name the nonmetal with an “ide” (cation then anion)
 All ionic compounds since they have a metal and a nonmetal, are salts
Ex: NaCl (Sodium Chloride), MgBr2 (Magnesium Bromide)
 Use stock system when there is more than one oxidation number
Ex: PbCl4 – Lead (IV) Chloride
 The old naming has the first oxidation as “ous” and the second as “ic”
Ex: FeCl3 (Ferric Chloride), FeCl2 (Ferrous Chloride)
st
 1 ox. names: Iron (Ferrous), Copper (Cuprous), Tin (Stannous), Gold (Aurous), Lead (Plumbous), Cobalt (Cobaltous),
Manganese (Managnous), Mercury (Mercurous), Nickel (Nickelous), Chromium (Chromous)
 The formula for mercurous/mercury(I) is Hg2+2
Ternary Compounds:
 The polyatomic ion (radical) is named with “hypo --- ite”, “ite”, “ate”, “per---ate”
 Atoms in the same group have the same pattern as another element in the same group
Ex: NaClO3 (Sodium chlorate) and NaBrO3 (Sodium bromate)
 Use stock system if necessary
Ex: Fe(NO3)2 – Iron (II) Nitrate or Ferrous Nitrate
Binary Molecular Compounds:
 Use prefixes, but don’t use mono on the first element
 The second element also ends in “ide”
Ex: N2O3 (Dinitrogen trioxide)
 If there are two conflicting vowels right next to each other, one of them is dropped
Ex: CO (Carbon monoxide)
Binary Acids: * Acids are added to water, not the other way around
 Name the acid as “Hydro---ic” acid
Ex: HCl (Hydrochloric acid)
Ternary Acids:
 Take the polyatomic name as a ternary compound and turn “ate”  “ic” and “ite”  “ous”
Ex: H3PO4 – Phosphoric acid
 HCN is an exception to the rule: Since it doesn't have an "ate" or "ite," you name it with the binary method
Ex: HCN – Hydrocyanic acid