Chem Exam Tomorrow And Haven't Started Revising

[RedRival]

Originally posted by Spider_dude@23 November 2003 - 13:37
what are the main topics and i can give you a quick summary about them.

redox, hydrocarbons, gas laws, energy profile diagrams, metal displacement reactions, etc...wow I just came up with those on of the top of my head.

[3rd gen noob]

with spider helping, you should do well

[pol]

Originally posted by 3rd gen noob@23 November 2003 - 20:59
with spider helping, you should do well 

[internet.news]

in chemistry there are some stuff you can summarize
you have to imagine some special reactions

~nice dreams...

[kAb]

here are my chem notes from 3 different pages (we are allowed to have 1 page of notes with the test)

(btw this is only like early chapters)

Laws of conservation mass – In products of a chemical reaction have the same mass as the reactants
Law of conservation energy – Any physical or chemical change, energy can change form, but toal amount of energy remains constant. Cannot be created or destroyed.
Specific Heat capacity – amount of heat energy required to raise the temp. 1g of the substance by 1k. SI unit for energy is joule.
Accuracy-how exact it is. How close to true value. Precision how closely several measurements of the same quantity made in the same way agree w/ 1 another.
Cathode Ray Tube Experiment – J.J Thompson – Pumped most of the air out of the glass tube, Applied a voltage to two metal plates placed at each end of the tube. The plates are called Electrodes. One of the plates is called the anode, attached to positive terminal of the voltage source, therefore it had a positive charge. The other plate, a cathode carried a negative charge, because it was attached to the negative terminal of the voltage source.
Rays emerged from cathode and struck the anode & nearby glass walls. Cathode rays are called that because of their origin. Used in T.V., monitors, oscilliopes, & radar displays.
Cathode rays consist of tiny particles, they have mass. Negatively charged. Electron mass is 5.485799x10 to the negative 4. Cathode rays always identical.
Rutherford Experiment – A beam of positively charged particles called alpha particles directed at gold foil. Most particles shot at foil and passed through. But some were deflected. Only a very concentrated positive charge located in the gold atom could repel the fast moving positively charged alpha particles. Came up with the nucleus. It meant that most of the atom was made of empty space though. Nucleus contains all of atom’s positive charge and nearly all its mass, but only a small fraction of its volume. Nucleus less than 1/10,000 of the radius of the whole atom.
Protons – Positively charged particles.
Protons & Neutrons found inside nucleus. Electrons found outside nucleus.
Atomic number-number of protons. # of protons = # of electrons.
Mass number minus Atomic number = number of neutrons
Isotopes-atoms of same element w/ diff #’s of neutrons. One of two or more atoms having the same atomic number but different mass numbers.
Radioisotopes – unstable nuclear configurations.
Allotrope-one of a number of different molecular forms of an element.
Density=Mass/Volume
Organic compounds – have carbon
Electromagnetic spectrum
Heterogeneous mixture – Not evenly distributed (salad)
Homogeneous mixture – uniformly distributed (gas, syrup, air)
10 mm = 1cm. 100 cm=1m. 1000m=1 km
ml=cm cubed. Liter = 1000 milliliters. Liter = 1000cm cubed.
States of matter- solid liquid gas. Solid, particles packed together most closely, but they move.
Line spectrum. Light is emitted when excited electrons go down the ground state.
Electromagnetic spectrum – includes x-rays, ultraviolet and infrared radiation, microwaves, and radio waves. light=energy
Frequency X wavelength = speed of light
Element can be excited if electron is not where it is expected to be
Each orbital holds 2 electrons. 2, 8, 18, 32
S-spherical, P – dumbbell, D - 4 leaf clover. Shapes represent clouds.
Difference between 3p & 2p orbitals, 3p takes up more space.
2n squared to find electron number, where N is the principle energy level.
Sig Figs – addition, no more #’s than the least # of sig figs to the right of the decimal.
A number that can only take definite values is a quantum number.
Columns in periodic table called groups.
Metals are members of groups 1 – 12 as well as some elements in group 13 and 16. groups 3-12 = transition metals.
Nonmetals = second region of the periodic table.
Semiconductors or metalloids conduct electricity better than nonmetals but not better than metals.
Groups 1, 2, and 13-18 are called main group elements.
Alkali metals – react w/ water to produce alkaline solutions.
Halogens – combine w/ most metals to produce salts. Most reactive group of nonmetals.
Noble Gases – low reactivity
Hydrogen its own class- behaves unlike other elements because it has one proton and one electron.
Conversions: 1L/1,000ml 1,000g/1Kg

Rutherford Experiment – A beam of positively charged particles called alpha particles directed at gold foil. Most particles shot at foil and passed through. But some were deflected. Only a very concentrated positive charge located in the gold atom could repel the fast moving positively charged alpha particles. Came up with the nucleus. It meant that most of the atom was made of empty space though. Nucleus contains all of atom’s positive charge and nearly all its mass, but only a small fraction of its volume. Nucleus less than 1/10,000 of the radius of the whole atom.
Staircase, Elements to the left are metals, Elements to the right are nonmetals
Alkali Metals – Highly reactive metallic elements in Group 1 that react rapidly with water to form hydrogen and alkaline solutions that burn in air
Alkaline-earth metals – Reactive metallic elements with 2 electrons in the outermost energy level
Halogens – A non metallic element that has seven electrons in the outermost energy level that combines with many metals to form salts. Most reactive group of non metals.
Noble Gases – Low reactivity
Valence Electrons – Electron in the outermost energy level of an atom, where it can participate in bonding
Protons – Positively charged particles.
Protons & Neutrons found inside nucleus. Electrons found outside nucleus.
Atomic number-number of protons. # of protons = # of electrons.
Mass number minus Atomic number = number of neutrons
Isotopes-atoms of same element w/ diff #’s of neutrons. One of two or more atoms having the same atomic number but different mass numbers.
Density=Mass/Volume
Heterogeneous mixture – Not evenly distributed (salad)
Homogeneous mixture – uniformly distributed (gas, syrup, air)
10 mm = 1cm. 100 cm=1m. 1000m=1 km
ml=cm cubed. Liter = 1000 milliliters. Liter = 1000cm cubed.
Line spectrum. Light is emitted when excited electrons go down the ground state.
Electromagnetic spectrum – includes x-rays, ultraviolet and infrared radiation, microwaves, and radio waves. light=energy
Frequency times wavelength = speed of light
Element can be excited if electron is not where it is expected to be
Each orbital holds 2 electrons. 2, 8, 18, 32
S-spherical, P – dumbbell, D - 4 leaf clover. Shapes represent clouds.
Difference between 3p & 2p orbitals, 3p takes up more space.
2n squared to find electron number, where N is the principle energy level.
Sig Figs – addition, no more #’s than the least # of sig figs to the right of the decimal.
A number that can only take definite values is a quantum number.
Semiconductors or metalloids conduct electricity better than nonmetals but not better than metals.
Groups 1, 2, and 13-18 are called main group elements.
Hydrogen its own class- behaves unlike other elements because it has one proton and one electron.
Ion-An atom or group of atoms with an electrical charge
Sublevel= 5 orbitals
3rd main energy level = 9
Trends:
Across Periodic Table (Down Periodic Table)
Electronegativity Increases (Decreases)
Ionization Increases (Decreases)
Electron affinity increases (Decreases)
Shielding effect is constant (Increases)
Atomic Radius decreases (Increases)
OXIDATION:
0 unless otherwise stated
hydrogen = +1, combined w/ metal = -1
Fluorine = -1
Oxygen w/ compounds = -2 w/ fluorine = +2
Oxidation numbers added up must be zero
Across Periodic Table (Down Periodic Table)
Electronegativity Increases (Decreases)
Ionization Increases (Decreases)
Electron affinity increases (Decreases)
Shielding effect is constant (Increases)
Atomic Radius decreases (Increases)
1000ml/1L – 1000g/1kg
Density = mass/volume
Celsius to Kelvin: -273 C/0K
Sig Figs – addition, no more #’s than the least # of sig figs to the right of the decimal.
Electronegativity: Tendency of an atom to attract bonding electrons to itself when it bonds w/ another atom.
Cations-Ions with a positive charge
Anions – ions with a negative charge
OXIDATION:
0 unless otherwise stated
hydrogen = +1, combined w/ metal = -1
Fluorine = -1
Oxygen w/ compounds = -2 w/ fluorine = +2
Oxidation numbers added up must be zero
Oxidadtion # = charge on ion
In compounds elements of group 1, 2 and aluminum have positive oxidation numbers of +1, +2, +3 respectively.
Ionic compound – composed of oppositly charged ions
Subscripts – used to denote the number of atoms in a formula – balance charges
Ionization energy is required to remove an electron
Valence electrons- S & P
Nonmetals have more valence electrons
Octet rule – tendency of atoms of elements to gain or lose electrons so that their outer s and p orbitals are full with eight electrons
Atoms react to exist @ lower energy
Orbital-a region of an atom in which there is a high probability of finding on or more electrons
Sodium chloride – NaCl
Each orbital holds 2 electrons. 2, 8, 18, 32
Mass x 1mol/atomic mass x 6.02x10 to the 23
Or
Mass x 6.02x10|23/atomic mass
Elements that border zig-zag line are metalloids
Nonmetals in solid state are brittle
Energy level = # of possible orbital shapes
Ion-An atom or group of atoms with an electrical charge
Line spectrum. Light is emitted when excited electrons go down the ground state.
Electromagnetic spectrum – includes x-rays, ultraviolet and infrared radiation, microwaves, and radio waves. light=energy
Frequency times wavelength = speed of light
Alkali Metals – Highly reactive metallic elements in Group 1 that react rapidly with water to form hydrogen and alkaline solutions that burn in air
Alkaline-earth metals – Reactive metallic elements with 2 electrons in the outermost energy level
Halogens – A non metallic element that has seven electrons in the outermost energy level that combines with many metals to form salts. Most reactive group of non metals.
Noble Gases – Low reactivity
Valence Electrons – Electron in the outermost energy level of an atom, where it can participate in bonding


they repeat because i took 4 tests with 4 different pages of notes, but i took some notes from previous chapters

hope i could help

btw: this is HS chemistry, not college

[Spider_dude]

i can photocopy my university books for you, but it will take me about 3weeks.

[kurtsl0an]

hey wait a sec, this thread was created on a saturday - u had a chem exam on sunday?

[RedRival]

nah it's created on Sunday. and thanks guys for the notes, but i'm not even allowed notes here and some stuff on the previous page is a bit irrelevant, but still thanks heaps.

I just hope I don't screw up Gas laws and redox....or electorn dot diagrams...i suck at those and yehr i'm in high school in aus not college...

btw, as im posting this, i'm only an hr 15mins away from my exam.

[3rd gen noob]

Originally posted by RedRival@24 November 2003 - 03:15
nah it's created on Sunday. and thanks guys for the notes, but i'm not even allowed notes here and some stuff on the previous page is a bit irrelevant, but still thanks heaps.

I just hope I don't screw up Gas laws and redox....or electorn dot diagrams...i suck at those and yehr i'm in high school in aus not college...

btw, as im posting this, i'm only an hr 15mins away from my exam.

any last minute questions you need help with?

[kAb]

how do you felt u did?