Honors Chemistry Semester 1 Review

Chapter 1

Mass – amount of matter

Volume – amount of space an object occupies

Plasma – 4^th state of matter, ionized gas

Atom – smallest unit of matter

Element – a pure substance made of one type of atom that cannot be broken down to simpler, stable substances

Compound – a substance made up of two or more elements chemically bonded together

Molecule – a group of atoms bonded together, representing the smallest unit that can take part in a chemical reaction

Diatomic element – two atoms of the same element bonded together

Mixture – a blend of two or more kinds of matter, each of which retains its own identity and properties

Homogenous mixture – a mixture in which the composition is the same throughout (ex. Solution)

Heterogeneous mixture – a mixture in which the composition is not the same throughout (ex. Milk, blood, trail mix)

Extensive property – does depends on the amount of matter

Extensive property examples – mass, volume, amount of energy

Intensive property – does not depend on the amount of matter

Intensive property examples – density, b.p., m.p., conductivity

Metals – solids at room temp, luster, malleable, ductile, conductors

Nonmetals – liquids and gases at room temp, brittle, poor conductors

Metalloids – solids at room temp, luster, semi-conductors

Distillation – separating liquids based on boiling point

Decanting – separating a solid and a liquid by pouring off the liquid after the solid settles

Filtering – separating a solid and a liquid by pouring the mixture into a filter – the liquid will pass through but the solid will remain

Evaporation – a method to separate water from a dissolved liquid

Chromatography – passing a mixture a medium where the compounds move at different rates

Chapter 2

Accuracy – closeness of a measurement to the accepted value

Precision – closeness of a group of measurements

Percent error – used to determine accuracy

Percent error formula – [(experimental – accepted) / accepted] x 100

Significant figures – all the digits known with certainty plus an estimated digit

Sig Fig Rules –

1.      All nonzero integers are significant

2.      Captive zeroes are always significant

3.      Leading zeroes are never significant

4.      Trailing zeroes are significant if there is a decimal point

add/sub whole number sig fig rule – the final sig fig is in the same place value as the left most uncertain digit

add/sub decimal sig fig rule – as many decimal places as the least precise piece of date

multiply/divide sig rule – as many total sig figs as the least precise piece of data

Scientific method – a logical approach to solving a problem

Quantitative data – numerical data

Qualitative data – observations made uses senses

Control – the standard used for comparison in an experiment

Theory – hypothesis that has withstood many tests

Tera – T, 10^12

Giga – G, 10^9

Mega – M, 10^6

Kilo – k, 10^3

Hecto – h, 10^2

Deca – da, 10^1

Base – B, 10^0

Deci – d, 10^-1

Centi – c, 10^-2

Milli – m, 10^-3

Micro - µ, 10^-6

Nano – n, 10^-9

Pica – p, 10^-12

Dimensional analysis – mathematical technique that allows one to use units to solve problems

Derived units – a combination of units, used as a conversion factor

Density – ration of mass to volume (m/V)

Chapter 3

Law of Conservation of Mass – LAVOISIER: mass is neither created nor destroyed during ordinary chemical reactions or physical changes

Law of Definite Proportions – PROUST: A single chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound

Law of Multiple Proportions – DALTON: if two or more compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers

Dalton’s Atomic Theory – (* revised)

1.      All matter is composed of small particles called atoms

2.      Atoms of an element are identical in size, mass, and other properties*

3.      Atoms cannot be subdivided*, created, or destroyed

4.      Atoms of different elements combine in simple whole number ratios to form compounds

5.      in chemistry ratios, atoms are combined, separated, or rearranged

Thompson’s Atomic Theory – concluded that negatively charged particles called electrons existed – atoms are electrically neutral, so atoms must contain a positive charge to balance the negative electrons

Cathode ray tube experiment – THOMPSON: rays were deflected away from negatively charged objects

Millikan’s conclusions – used the charge to mass ratio to determine the mass of an electron (so small it’s considered negligible), concluded that other particles must be present to account for the mass

Oil drop experiment – MILLIKAN: measured the charge of an electron

Rutherford’s conclusions – the alpha particles hit something small and dense, called a nucleus

Gold Foil experiment – 1 in 8000 alpha particles were deflected back

Isotopes of hydrogen – protium (no neutron), deuterium (1 neutron), tritium (2 neutrons)

Alpha emission – least penetrating, blocked by a sheet of paper

Alpha particle – ⁴₂He

Beta emission – more penetrative, blocked by aluminum foil

Beta symbol – ⁰_-1e- OR ⁰_-1β

Positron emission – more penetrative, stopped by aluminum foil

Positron symbol – ⁰₊₁e- OR ⁰₊₁β

Gamma emission – electromagnetic radiation, stopped by lead

Gamma symbol - ⁰₀Ƴ

Symbol for an isotope – hyphen notation (ex. Lithium – 7)

Percent abundance ex. – Lithium – 6. (2/32) x 100

Average atomic mass – (mass x % abundance) + (mass x % abundance)

Mass used in mole-gram conversion – atomic mass of the element

Avogadro’s number – 6.022 x 10^23

Chapter 4

Wavelength and frequency – inversely proportional

Frequency and energy – directly proportional

Photoelectric effect – light that strikes the surface of a metal causes the metal to emit an electron. The light must have a specific frequency depending on how tightly the electrons are bonded to the metal

Speed of light – 3.00 x 10⁸ m/s

Frequency (f) x wavelength (λ) – speed of light (c)

Plank’s constant (h) – 6.626 x 10^-34 Js

Plank’s constant (h) x frequency (f) – energy

Given wavelength (λ), want Energy (Jules) – E = (hc) / λ

Emission-line spectrum (H) – add energy, light is passed through a prism (red, blue/green, blue, violet)

Excited state – when an atom absorbs energy, its electrons move to a higher energy level

Ground state - The lowest energy state of an atom, when electron goes from the excited to ground state, light is produced

Chapter 4

Louis de Broglie’s research – Treats e- as waves with a certain probability of being found at different distances from the nucleus

Werner Heisenberg Principle – It is impossible to determine simultaneously both the position and velocity of an electron

Erwin Schrodinger’s research – Treated e- as waves having a certain probability of being found in orbitals at various distances from the nucleus

Orbital – A three-dimensional region around the nucleus that indicates the probable location of an e-

Quantum Number 1 – Principal Quantum Number

Principle Quantum Number – Indicates the main energy level around the nucleus, symbol = n

Quantum Number 2 – Angular Momentum Quantum Number

Angular Momentum Quantum Number – Tells the shape of an orbital (sublevels)

S orbitals – Spherical shape, lowest E

P orbitals – Dumbbell shape, low E

D orbitals – Double dumbbell shape, high E

F orbitals – Flower shape, highest E

How many orbitals in S sublevels – 1 orbital

How many orbitals in p sublevels – 3 orbitals

How many orbitals in d sublevels – 5 orbitals

How many orbitals in f sublevel – 7 orbitals

Quantum Number 3 – Magnetic Quantum Number

Magnetic Quantum Number – Indicates the orientation in space of orbital (x, y, z axis), symbol = m

Quantum Number 4 – Spin Quantum Number

Spin Quantum Number – Two e- can exist in one orbital but must have opposite spin states (+1/2; -1/2)

Huns Rule – Each orbital receives 1 e- before receiving a second e-

aufbau Principle – e- will occupy the lowest energy position (must fill up 1s before 2s, etc)

Pauli Exclusion Principle – no two e- can have the same Spin State

Valence e- – e- in the highest main energy level

Chapter 5

Group 1 – alkaline metals

Alkaline metal properties – shiny, soft, most reactive metals, not found as free elements in nature and must be stored in oil

Group 2 – alkaline earth metals

Alkaline earth metal properties – not found as free elements in nature, not as reactive as group 1, harder, denser, and stronger than group 1, higher melting point that group 1

Group 13 properties – metals, found in nature as compounds, harder and denser than group one elements but softer and less dense than transition metals

Group 17 – halogens

Halogen properties – most reactive nonmetals, bonds to itself

Group 14 – noble gases

Noble gas properties – unreactive, complete valence shell

Cation – positive ion, formed by losing electrons

Anion – negative ion, formed by gaining electrons

Cation, anion, and atom size – cations -> atoms -> anions

Electron shielding – more electrons lie between the nucleus and the electrons in higher energy levels

Group vs period electrons – Group: electrons occupy sublevels in higher energy levels (higher distance between nucleus and valence e-, lower force of attraction). Period: electrons are added to the same energy level (e- shielding remains the same, and no significant change in atom size)

Chapter 6

Ionic bond – a metal transfers electrons to a nonmetal

Nonpolar covalent bond – nonmetals share electrons equally

Polar covalent bond – nonmetals share electrons unequally

Ionic bond electronegativity – 1.8 or greater

Polar covalent bond electronegativity – 0.4 to 1.7

Nonpolar covalent bond electronegativity – 0.0 to 0.3

Ionic compound – compound of positively and negatively charged ions combine for a net charge of 0

Covalent compound properties – molecule structure and formula

Ionic compound properties – crystal lattice structure, formula unit

Metallic bond properties – crystal lattice structure, one type of metal atom