Chem Exam 1
Metric Conversion |
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Significant Figures |
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Chapter 1.1 Chemistry Mass vs Matter Building Blocks of Matter Properties & Changes in Matter Kinetic Theory of Matter |
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Matter 1.2 Physical Property Chemical Property Physical Change Chemical Changes Classification of Matter Matter Flowchart Pure Substances Element Compound Pure Substances Mixtures | Can be observed without changing the identity of the substance
Describes the ability of a substance to undergo identity changes
Changes the form of substance without changing the identity and the properties will remain the same
Changes the identity of the substance and the products will have different properties
Some signs include:
It’s one of those yes - no charts, examples: Graphite - Element Pepper - Heterogeneous Mixture Sugar (sucrose) - Compound Paint - Heterogeneous Mixture Soda - Solution, or Homogeneous Mixture Composed of identical atoms, examples: copper wire and aluminum foil Composed of two or more elements in a fixed ratio, and it has properties different from those of individual elements, like table salt (NaCl) Pure substance always contains the same, fixed ratio of elements Pure substance has exactly the same characteristic properties Ex: C+O = CO and C+O+O = CO2, both have definite and different compositions Variable combination of two or more pure substances which retain their own properties. There’s types Solutions:
Colloid:
Suspension:
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Intro to the Periodic Table 1.3 Atomic Numbers Atomic Mass Electrons, Protons & Neutrons Organisation History | The Atomic Numbers show the number of protons, usually found in the left corner or bottom left when specifying Average Mass of one atom, usually found below atomic symbol or top left when specifying
so atomic mass isn’t always the same in each version
The number of electrons is always equal to the protons. To find neutrons, atomic mass is just AM-P=N Electrons will equal protons unless it is charged. Groups, or Columns, have similar chemical properties Periods, or Rows, do not necessarily have similar chemical or physical properties Dimitri Mendeleev made the first table in 1896. He’s Russian. He organised elements by increasing atomic mass, elements with similar properties were grouped together. There were some discrepancies. Henry Moseley, a British man, in 1913, organised elements by increasing atomic numbers and that resolved the discrepancies. This would be the model used today. |
Chp 2.1 Scientific Method Model Theory Law SI Base Units Length Mass Time Temperature Amount of substance SI derived units Area Volume Density Energy Density Conversion factors |
Explanation as to how phenomena occur or how data or events are related (Ex: Bohr atomic model) Broad generalisation that explains a body of facts or phenomena (Ex: Big Bang Theory) Proven scientific fact, sometimes a math equation (Ex: E = MC^2) l, uses metres m, uses grams t, uses seconds T, uses kelvin (K) n, uses mole (mol) A, square metre V, cubic metre or litres D, Kilograms per cubic metre E, joule (J) Density = mass / Volume Ratio derived from equality between two different units that can be used to convert one to the other. It can be found if you know the relationship between the two units (ex: 100 cm to 1 m or 1 m to 100 cm) |
Chp 2.2 Accuracy Precision Percent Error SigFigs Again Scientific Notation Direct variation Indirect variation | How close a measurement is to the accepted value (correct) How close a series of measurements are to each other (consistent) Indicates the accuracy of a measurement Formula: Error = (Accepted value - Expected value) / Accepted value X 100 The last digit is called the Digit of Uncertainty. Move decimal until there’s one digit to its left, places moved = exponent Large # (1 >) → positive exponent Small # (< 1) → negative exponent Y = x Y = 1/x |
Chp 3.1 Law of Conservation of Mass Law of Definite Proportions Law of Multiple Proportions Dalton's Atomic Theory Subatomic particles Masses of Atoms Mass Number Isotopes Relative Atomic Mass Average Atomic Mass | Mass is neither created nor destroyed during ordinary chemical reactions or physical changes
A given chemical compound always contains the same elements in the exact same proportions by mass
If two different compounds are composed of the same two elements, then the ratio of the masses of the second element are small whole numbers Proposed as an explanation for the laws in 1808, when he reasoned that all elements were composed of atoms and only whole numbers of atoms can form compounds Atoms → Electrons and Nucleus → Electrons → Negative charge & Nucleus → Protons and Neutrons → Protons → Positive charge & Neutrons → Neutral charge Mass # = protons + neutrons
Atoms of the same element with different mass numbers
amu - atomic mass unit Based upon carbon Simplified system for masses of elements
Weighted average of all isotopes
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Chp 3.2 Mole Molar Mass STP | It’s very large A counting number (like a dozen) Avogrado’s number - (N^A) 1 mol = 6.02 * 10^23 items Mass of 1 mole of a pure substance
Usually rounded 2 decimal places Molar masses of elements are found on the periodic table as the atomic mass For a compound: Molar Mass equals sum of molar masses of each individual element Grams (divide by molar mass) → Mole (multiply by avogardo’s number)→ Atoms Atoms (divide by avogardo’s number)→ Mole (multiply by molar mass) → Grams Never abbreviate molecules Standard Temperature and Pressure Temperature = 25 C = 22.4 L |
Chp 7 Percent Composition Empirical Formula Molecular Formula | = mass of element (molar x how many times it is present) / total mass * 100 You can multiply the percent of an element by the total mass to get the element’s mass Smallest whole number ratio of atoms in a compound
The steps:
The “True formula” - actual number of atoms in a compound The Steps:
Empirical formula mass is the molecular mass of the empirical formula. |
Chp 4 Wavelength Frequency Amplitude Long wavelength and low frequency Short wavelength and high frequency Frequency and Wavelength Quantum Theory Quantum Theory - energy of a photon formula | Length of one complete wave, upside down Y (lamda)
# of waves that pass a point during a certain time period, cycles per second
Distance from the origin to the trough or crest
Means low energy Means high energy Inversely proportional, c = V C = speed of light (3.00 * 10^8 m/s) = wavelength (m, nm, etc, m = 1*10^9 nm) V = frequency (Hz) Planck (1900)
Einstein (1905)
Energy of a photon is proportional to its frequency: E = hV E: Energy (J, joules) h: Planck’s constant (6.6262 10^-34 J s) V: frequency (Hz) |
Chp 4.2 Bohr Model | In atoms, there are energy levels at different distances with different numbers of electrons
Level – Electrons N = 1 – 2 N = 2 – 8 N = 3 – 18 N = 4 – 32 N = 5 – 32 N = 6 – 18 N = 7 – 8 |
Chp 4.3 Quantum Model Heisenberg Uncertainty Principle Four Quantum numbers Pauli Exclusion Principle | Louis de Broglie (1924)
Impossible to know both the velocity and position of an electron at the same time
Specify the address of each electron in an atom
Orbitals combine to form a sphere shape
No two electrons share the same ‘address’, or four quantum numbers Each one has a unique address
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Chp 4.4 Aufbau Principle Hund’s Rule Noble Gas Configuration Stability | Electrons fill the lowest energy orbitals first, seeking stability and low energy, or “Lazy tenant rule”
Neon = 1s^2 2s^2 2p^6 Oxygen = 1s^2 2s^2 2p^4 powers represent electrons in the orbital You can tell based on columns in those blocks Within a sublevel place one electron per orbital before pairing (empty bus rule) You fill each empty orbital with 1 electron before pairing electron You go to the most recent noble gas before it
As [Ar] 4s^2 3d^10 4p^3 When an orbital is half filled, it is more stable than something partially filled |
Balancing Equations | Write equation Count atoms on each side Find the ratio and what number is needed as the coefficient to balance it If an element appears more than once per side balance it last |
Mole to Mole Ratios | Moles are represented by subscripts in a compound & the coefficients in a balanced chemical equation |