Craddock is phat
sup papa craddock, homie rawr, i roast you and eat you. and evolve rawr rawr rawr lul
i'm not a grammar expert, but who wrote this? My half-lucidness can't study. So horrible... -HS
Chapter 1

1. kinetic molecular theory of matter: help to explain the forces between molecules in solids, liquids, and gases
   • in solids, particles are packed closely together in a pattern, causing them to vibrate as a single unit.
   • a special case is water in which the hydrogen bonds freeze. because hydrogen bonds require a certain distance to form, the water particles freeze farther apart, making it less dense than its liquid state(chomander was here)
     
   • in liquids, the atoms are arranged randomly (not in a solid pattern), which explains their fluid-like characteristic
   • in gases, particles are very far apart and move very rapidly, colliding randomly with other atoms. The random motion of the gases results in the volume of the gas to take up the volume of the container.
2. According to kinetic molecular theory, the higher the temperature, the faster the particles move.
   
   • the kinetic energy of a substance determines the state of the substance
     • if the particles of a solid are heated enough, their kinetic energy causes the breaking of bonds and results in a liquid
       
3. Matter can be classified as homogeneous or heterogeneous matter.
   
   • Heterogeneous matter: nonuniform composition that can be physically separated into homogeneous matter
     
   • Homogeneous matter: uniform composition
     
4. Mixtures: homogeneous v. heterogeneous
   • heterogeneous mixtures: nonuniform composition even if it may appear to be otherwise; the composition of one area differ from those in another
     • examples: blood, milk, soup
   • homogeneous mixtures: AKA Solutions: two or more substances in the same state; properties uniform throughout substance
     • examples: air, salt water, gasoline
       
   • mixtures can be physically separated into pure substances through methods such as filtering, sublimation, extraction, and distillation.
     • Pure Substances: substances that have a unique set of physical and chemical properties and cannot be separated into different substances by physical means
5. Element: a substance composed of only one type of atom
   
   • organized by atomic number on the periodic table
     
   • 90 elements found in nature
6. The atom is the smallest particle of an element that retains all the properties of the element.
   
   • atoms contain protons, neutrons, and electrons
   • the number of protons (atomic number) determines the type of element and its placement on the periodic table
7. Chemical bond v. Chemical compound
   • chemical bond: a bond that holds together two elements
     
     • water, salt, and sugar have chemical bonds but are chemical compounds
       
   • chemical compounds: pure substances that are composed of two different elements
     
     • when elements form bonds and become a compound, their original properties are changed
       • original color, hardness, boiling point, and other properties evolve into properties of the compound
     • water, sugar, and salt are compounds
       
   • unlike mixtures of elements, chemical compounds have a definite percentage composition (by mass) of their elements. Plus the properties of chemical compounds differ from the properties of the individual elements.
     
8. Physical properties are those that can be observed and measured without chemical reactions or changing the composition of the substance.
   • examples: color, state, melting/boiling point, density, solubility, electric conductivity, malleability, ductility, viscosity
     
   • Density = mass / volume
     
9. Temperature influences the numerical value of certain properties, including the density of a substance.
   • for water: the volume of a the liquid is smaller than the solid (ice is more less dense than liquid water)
   • To change from Celsius to Kelvin, add 273.15K
     
10. Extensive v. Intensive Properties:
    
    • extensive: dependent of the quantity of the substance
      
      • mass, volume
        
    • intensive: not dependent on the quantity of the substance
      
      • melting/boiling point, density
        
11. Changes to the physical properties of a substance are physical changes.
    • the physical identity of the substance does not change despite changes to size and shape
      • state changes, shape changes

12. Changes to the chemical properties are chemical changes.

• substances undergo a chemical reaction during which one ore more substances ( reactants ) are transformed into one or more different substances (products)
  
• depicted by a chemical equation
  • left = reactants right = products

13. Qualitative vs. Quantitative observation

• qualitative observation-- non-numerical experimental observations, such as descriptive or experimental data. Example: color, appearance
• quantitative observation--Numerical experimental data such as measurements in mass or volume
• to convert from Fahrenheit to Celsius: T(°C) = (5°C)/(9°F)[T(°F)-32]

14. Precision vs. Accuracy

• precision - how close together measurements are
• accuracy - how close the measurements are to the accepted value
• percent error = (100%)(|theoretical - actual|)/(actual)
  

15. Guidelines for Determining Significant Figures:

1) To determine the number or significant figures in a measurement, read the number from left to right and count all digits, starting with the first digit that is not zero only with decimal point
2) When adding or subtracting numbers, the number of decimal places in the answer should be equal to the number of decimal places in the number with the fewest digits
3) In multiplication and division, the number of significant figures in the answer should be the same as that in the quantity with the fewest significant figures

Chapter 2

Electricity is the flow of electrons. There are two types of electric charge - positive and negative. Charge builds up when one substance rubs over another, which implies substances are charged when separated.

• Opposites attract.
• Likes repel.
• Equal amounts of opposite charges cancel out.
  

No one cares about radioactivity T___________T;;

Dalton's Atomic Theory

1. Each element is made up of tiny, indestructible particles called atoms.
2. All atoms of a given element has the same mass and other properties that distinguish them from other elements.
3. Atoms combine in whole number ratios to form compounds.
4. In a chemical reaction, atoms of one element cannot change into atoms of another element. They only rearrange themselves.

Thompson's Cathode Ray Tube

A Cathode Tube is composed of a Cathode (negative) and a Anode (positive). Thompson shot streams composed of tiny particles moving in one direction from the cathode to the anode. Near the anode, Thompon placed a positive and negative plate. He observed the particles being deflected towards the positively charged plate. Then, he added magnetic fields perpendicular to the tube and countered the positively charged plate's effect. With this, he was able to calculate the charge/mass ratio of the "electron", which was -1.76 x 10^8 Coulombs/g.

Miliken's Oil Drop Experiment

Oil was sprayed as droplets into a container. TH\he droplets of oil collect electrons as they fell. A negatively charged plate was placed on the bottom of the container. With a microscope, Miliken observed the droplets of oil were floating. By varying charge on the plate and recording different heights of droplets, in addition to Thompson's charge/mass ratio of the electron, Miliken calculated the charge of an electron to be -1.60 x 10^-15 Coulombs.

Plum Pudding Model

Thompson believed the atom was a sphere of positively charged "stuff" with negatively charged particles spread evenly throughout.

Rutherford's Gold Foil Experiment - Modern Atomic Theory

Rutherford believed that if he shot an alpha particle through a sheet of thin gold metal, the particles would go straight though. (Like throwing a rock through tissue paper). However, what he observed was that some were deflected. This was quite shocking to him (imagine that rock bouncing back at you...) and he concluded that atoms were composed of largely empty space, but with a dense center. He said that the atom has a very dense postively charged center (nucleus).

If you don't know what the atomic number of an element is....then...(fill in blank)

Atomic particles are expressed in atomic mass units (amu).

• 1 amu = 1.661 x 10^-24 = 1/12 the mass of a carbon atom. (So C weighs 12.000 amu)
• Protons and neutrons weigh almost 1 amu. The mass of an electron is very insignificant compared to them (1/2000 of a proton-ish).

A =Mass number = number of protons + number of neutrons
Z = Atomic number
X = Element symbol

A

ZX

Isotopes of Hydrogen:

• Protium = "hydrogen" (no neutons)
  
• Deuterium = "D" = hydrogen-2
  
• Tritium = "T" = hydrogen-3
  

Natural Abundance is the amount of each isotope of an element present in nature.
Percent Abundance is almost the same thing; it's the amount of each isotope of an element present in a given condition.

Percent Abundance = (Number of atoms of a given isotope)/(Total number of atoms of all isotopes of that element) x 100%
We determine the masses of the isotopes and their percent abundances experimentally using a mass spectrometer. Read the book for details. pg 70 There's a nice picture.

Mass Defect is the mass "lost" from calculating the mass of an element by adding its constituent particles (protons, neutrons, electrons) because energy is required to form bonds.

Since energy = mass (E=mc^2), that's where the mass was "lost".

Atomic Weight is the average weight of all isotopes of an element in their specific proportions.

Atomic Weight = (% Abundance of Isotope 1)x(Mass of Isotope 1) + (% Abundance of Isotope 2)x(Mass of Isotope 2) +...

Fun fact: the term "mole" was introduced by Wilhelm Ostwald, who derived the term from a Latin word meaning "heap" or "pile"! How exciting.

The Periodic Table (From Left to Right)
IA: Alkali Metals - VERY reactive, solid at room temperature. Only found as compounds in nature.
IIA: Alkaline Earth Metals - Fairly reactive, only found as compounds in nature.
3B-2B: Transition Metals - Less reactive, but most found as compounds. Some are stable elements (Ag).
Inner Transition Metals (Lanthanides/Actinides) - Transition metal properties. Most are radioactive. 3A-6A: Metalloids/Metals/Nonmetals - Pretty random properties. Very abundant on Earth. pg. 82-86 for more
7A: Halogens - Very reactive nonmetals. Form salts with Alkali Metals.
8A: Noble Gases - Unreactive. Not abundant. Colorless.
Diatomic Molecules are atoms that occur naturally as dual atoms. They are H2, O2, N2, Cl2, Br2, I2, F2 (HONClBrIF). List of most abundant elements on Earth's crust:

1. Oxygen
2. Silicon
3. Aluminum
4. Iron
5. Calcium
6. Sodium
7. Magnesium
8. Potassium
9. Titanium
10. Hydrogen <---I was like wtf...

Chapter 3

1. Molecule: smallest substance into which a pure substance can be broken down into and still retain all chemical and physical properties of the pure substance

• composed of at least two atoms, from one or more element, held together with strong chemical bonds

2. When elements form compounds, the individual properties of each element changes into the properties of the element = difference between elements and compounds
Formulas:

• molecular formula: no structural information; only indicates the composition of the molecule and the proportion of each element (H2O)
• empirical formula: relative amounts of each element. (C6H12O6 would be CH2O)
• condensed formula: indicates how atoms are grouped; no direct structural info
  • a compound with the same molecular formula but different condensed formula have different chemical properties
• structural formula: high level of structural information; shows how atoms are attached within the molecule; line between atoms represent chemical bonds

3. Physical and chemical properties of a compound are defined by the structure of the molecule

• example: the structure of ice molecules account for ice being less dense than the liquid water because the molecules of liquid water are not packed as tightly as the ice molecules
  

4. Structures of molecules:

• ball and stick model: spheres represent the atoms and sticks represent the chemical bonds between the atoms

• space filling model: take into account the relative size of each atom and their proximities to each atom within the molecule; more realistic representations of molecules;
  

6. Ionic Compounds:usually a metal and nonmetal; composed of ions (atoms with either positive or negative charge)

• Cat (meow) ions: positively charged ion -- usually the metal
  
• anions: negatively charged ion -- usually the nonmetal
• monatomic ions: individual atoms that have lost or gained the election
  • metals of groups 1A-3A form cations with equal charge to the group number of the metal
    
  • many transition metals form 2+ and 3+ ions
    
  • nonmetals often form ions with a negative charge equal to 8 minus the group number of the element
    
  • hydrogen can either lose or gain an electron, depending on the other atoms in the compound

7. Properties of Ionic Compounds:

• positive and negative ions are attracted by electrostatic forces ("glue that holds ions of opposite charge together")
  
  • the force between the two ions is determined by Coulomb's law
    
    • remember physics: F=kQq / r^2 where k is a constant; Q is the charge of atom 1; q is charge of atom 2; r is distance between
• crystal lattice structure: 3D network arrangement of millions of positive and negative ions in ionic compound
  • such structure define set of physical and chemical properties of ionic compound
    • because each ion is surrounded by oppositely charged ion, the compound is very tightly bound
      
      • this tightness in the ionic compound requires lots of energy for the forces to be broken--reason for the very high melting points of ionic compounds
        
      • also reason why most ionic compounds are "hard solids"

8. Covalent Compounds: i.e. water, sugar; can be either solids, liquids, or gases
binary compounds: composed of two nonmetals

• hydrogen forms binary compounds with all nonmetals except noble gases
  
• most binary molecular compounds of nonmetals are a combination of one element from groups 4A-7A and hydrogen
• When naming binary compounds, the number of atoms of a given type in the compound is designated with a prefix, such as "mono-" "di-," "tri-," "tetra-," "penta-," "hexa-," and so on

9. Average Atomic Mass (amu) = Weighted average of all of the naturally occuring isotopes of a given element.

• Average Atomic mass = (% abundance of isotope 1 x mass of isotope 1) + (% abundance of isotope 2 x mass of isotope 2) + ...+
• 1.661 x 10^-24 g = 1 amu = 1/12 of Carbon-12

10. Percent Composition: any sample of pure compound consists of the same elements combined in the same proportion by mass. % of each element in a compound.
% of element = (Molar Mass of Element)/(Molar Mass of Compound)

• molecular composition can be expressed in 3 ways:
  1. in terms of number of atoms of each type of molecule through chemical formula of compound
  2. in terms of mass of each element per mole of compound
  3. in terms of the mass of each element in the compound relative to the total mass of the compound = mass percent

Random Laws:

• Law of Definite Proportions: All compounds have the same proportions of their constituent elements. - Proven by mass ratio.
• Law of Multiple Proportions: When two elements form two different compounds, the masses of one element will also be a whole number ratio. Ex. (XcompoundA (g))/(YcompoundB (g)) = whole #

Differences among hypotheses, laws, and theories:

• hypothesis- tentative explanation or prediction based on experimental observations
• law- concise verbal or mathematical statement of a behavior or a relation that seems always to be the same under the same conditions
• theory-unifying principle that explains a body of facts and laws based on them.

Formulas

• Density= Mass/Volume
• to convert from Fahrenheit to Celsius: T(°C) = (5°C)/(9°F)[T(°F)-32]
• percent error = (100%)(|theoretical - actual|)/(actual)
• Average Atomic mass = (% abundance of isotope 1 x mass of isotope 1) + (% abundance of isotope 2 x mass of isotope 2) + ...+
• 1.661 x 10^-24 g = 1 amu = 1/12 of Carbon-12
• Coulomb's law: F=kQq / r^2
• % of element = (Molar Mass of Element)/(Molar Mass of Compound)