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Summary ON Curriculum - Course Notes Gr 12 Chemistry

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Course
Chemistry

Institution
12th Grade

Covers in-depth complete course notes for grade 12 chemistry (university level). Includes topics: 1. Structure and Properties of Matter - History of the Atomic Theory - Various historic models and theories - Electromagnetic Spectrum - Schrodigner’s Equation - Quantum chemistry - VESPR ...

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Uploaded on September 30, 2023
Number of pages 36
Written in 2019/2020
Type Summary

electromagnetic spectrum
atomic models
schrodigners equation
quantum chemistry
vespr
thermodynamics
structure and properties of matter
history of the atomic theory
energy changes and rates of reaction

Institution Secondary school
Education 12th Grade
Course Chemistry
School year 4

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Chemistry Notes - Unit One: Structure and Properties of Matter

The History of the Atomic Theory
- Using a model
- Uses familiar ideas to explain unfamiliar facts observed in nature
- Can be changed as new information is collected
- We cannot see the inside of an atom, but use model to predict what’s happening
- Image, descriptions, mathematical…
- Timeline of atomic theory
- Started in 400 BC (Greece) - atoms looked like a billiard ball
- Modern science as been done in the last 200-400 years

Democritus (400 BCE)
- Greek philosopher searching for a description of matter
- What would happen if matter was divided into smaller and smaller pieces (is there a limit?)
- His theory
- Matter would not be divided into smaller piece forever, eventually the smallest
would occur
- Atomos: not to be cut
- Atoms were small, hard particles that were made of the same material but were
different shapes and sizes
- Aristotle, plato were more influential that him, and they believed that the four elements
were the approach to the nature of matter
- Democirtus’ theory was visited more in the 1800

Dalton’s Model (early 1800s)
- English chemist performed a number of experiments that led to the acceptance of atoms
- He was able to break down water and measured the mass ratio of H to O (8:1) and volume
ratio (2:1), without knowing about the atoms theory (evidence that atoms are real)
- His theory
- Took Democritus’s theory and combined it with his own discoveries
1. All matter is made of indivisible particles called atoms
2. Atoms of the same element are similar to each other (shape and mass)
3. Atoms of different elements differ from each other (shape and mass)
4. Atoms can neither be created nor destroyed
5. Atoms combine in whole number ratios to form compounds
- Foundation of modern chemistry (stoichiometry)

Thomson’s Model (1897)
- plum pudding model: atoms were made from positively charge substances with negatively
charged electrons scattered about
- Used a cathode ray tube where when negative particles hit the phosphorus screen, it would
light up, and applying an electric field the particles’ path would change
- magnetic fields also deflected the path as well
- Determine using the angles to find a mass to charge ratio

, - Experiments leading to the Thomson Model
- Current through gas
- Cathode rays coming from the path were negatively charged particles
- Conclusions
- Negative charges came from within the atom
- Atom is divisible (smaller particles had to exist)
- Corpuscles: negatively charged particles
- Gas was neutral so the atoms had to contain an equal amount postivice charges

Rutherford (1908)
- Gold foil experiment: fired positive particles (alpha particles) from a radioactive source at a
very thin sheet of metal foil
- Most went through and hit the fluorescent screen around it, but some got deflected
on the gold foil (sometimes at extreme angles)
- Since he thought that an atom had no overall charge, he thought the alpha
particles would go straight through
- The alpha particles hit “something solid” in the foil and deflected/bounced (a proton
in the nucleus of an atom in the foil)
- Conclusions
- Atoms are mostly empty space
- Atoms have a small dned positively charged center that repelled the alpha particles
- Rutherford’s “nucleus” and is tiny compared to the atom as a whole
- All of an atom’s positively charged particles contain nucleus
- Negatively charged particles scattered outside the nucleus around the atom’s edge

Bohr Model (1913)
- Problem with Rutherford’s model
- The negative and positive particles would be attracted to each other and collapse
- Electron will release energy while orbiting a nucleus
- Due to loss of energy, it would spiral towards the nucleus
- His model
- Electrons move in definite orbits, energy levels, or shells around the nucleus
- They are located at certain distances from the nucleus
- Apply a force to move it further from the nucleus and has more energy further away
from the nucleus (energy levels is the best term)
- The electrons can only be on the orbits, cannot exist in between (when it moves to a new
orbit, it just jumps there with no travel path and it need energy for it to jump)
- Give electron energy, it will jump to the next orbit (same when jumping down)
- Just the right amount of energy, just the right distance for the orbit: QUANTIZED
- Energies at specific wavelengths (kind of light that has different energy)
- It can also release same photon light (wavelength)
- An excited hydrogen atom, we would expect to see particular wavelengths of light produced
as the electrons fall back and emit photons of light corresponding to those energies
- Only particular wavelengths of light are emitted (wavelengths that correspond to the
required amount of energy to jump between levels)

,Who Said What?
indivisible electron nucleus orbit

Demo X

Dalton X

Thomson X

Rutherford X X

Bohr X X X

Textbook (pg. 173 #13): Quantized energy indicates quantized energy levels (electrons jumping from
one specific energy level to another corresponds to the energy of a specific colour of light).

Electromagnetic Spectrum
- Light exists in tiny packets called photons (properties of both waves and particles)
- Form of electromagnetic radiation, which can range from gamma to radiowaves
- The light we see is visible light (small portion of wavelengths) - 400 to 700 nm
- Animals can see more lights on the spectrum
- Different sizes of waves but all travel at the same speed
- Larger wavelengths (travel long distances, cell phones)
- Smaller wavelengths (help eliminate cancer cells or other medical uses)
- Wavelength: distance between two adjacent crests or troughs
- Frequency: number of waves per second
- ΔE=hf (energy of the light using Planck’s constant), c=λf (using speed of light)
- Waves on both ends of the spectrum (gamma - radio)
Short wavelength Long wavelength

High frequency Low frequency

High energy Low energy
- Using a prism we can diffract white light to see different colours/wavelengths
- Separates the light from longest to shortest wavelength
- Diffraction grating: hundreds of tiny lines etched into glass (or material)
- White light is a continuous spectrum (has all wavelengths of light within a range)
- No breaks between the different wavelengths
- Line spectrum has lights of light at specific and individual wavelengths
- Only able to see one wavelength at specific points on the spectrum
- Missing specific wavelengths
- Emission spectrum created from a hot material course (giving off light)
- Absorption spectrum is from a cold source (gas), blocks light from continuous source
- Absorption spectra: a source of white light is passed through a cold gas and a prism
- Taking away a select few wavelengths by absorbing them from the atoms in the gas
(different for each gas)

, - Emission spectra: a gas is given energy (by heat generally) until it glows
- The emitted light from the hot has is then passed through a prism (in a dark room)
- Only specific wavelengths are emitted by the atoms in the gas (coloured lines)

- High temperature, low pressure
- More electrons and more energy levels (then the more complicated the spectrum is)
- Hydrogen only has one energy level (very simple emission spectrum)
- Doesn't matter if emission or absorption spectrum, have lines in the same spots
- The difference in electron energy levels leads to the different colours (frequency)
- What causes colour
- Electrons only exist in energy levels
- Electrons in stable atoms and ions are in their ground state (lowest energy level)
- When energy is added to a element, the electron get excited from their ground state
to go to a higher energy level
- Then they will drop back down to lower energy levels, going back to their ground
state, and release extra energy in the form of light
- Therefore, as the gas release energy (release light) electrons come back down
- Exciting electrons
- n=1 is the ground state for an electron
- When energy is added, the electron raises to the next levels n=2, n=3
- Electron will fall back down and immediately
- When electron falls back to a lower energy level, the excess energy is emitted as
light (therefore we can see colour at specific wavelengths)
- Electrons emit energy when going down, absorbing energy when they move up
- There are only a specific and limited number of amounts of energy that electrons
can lose and gain
- When electrons are jumping down orbits, it’s showing an emission spectrum
- Can’t see some of the lines that are emitted since they’re outside of visible spectrum
- For emission/absorption: distance → energy → frequency → wavelength → colour
- For hydrogen spectrum: n=3-6 to n=2, is visible light (Balmer series), so we
can see it, but unable to see UV and infrared
- Pashen (infrared) is too small to be seen since it has a lower energy
therefore lower frequency, larger wavelength, so we can’t see (n=7)
- Lyman (UV): drop to n=1 from anywhere will result in energy release in the
form of UV radiation (undetectable) - since the distance was too large
- n=infinity, the electron has enough energy to leave the atom

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