SCH4U - Structure and Properties of Matter Flashcards
Electrons
An electron is a negatively charged subatomic particle that can be either bound to an atom or free (not bound). An electron that is bound to an atom is one of the three primary types of particles within the atom
Symbol: e^-1
mass: 9.109 x 10^-31
Protons
A proton is a subatomic particle found in the nucleus of every atom. The particle has a positive electrical charge, equal and opposite to that of the electron.
Symbol: P^1+
mass: 1.673 x 10^-27
Neutrons
A neutron is a subatomic particle found in the nucleus of every atom except that of simple hydrogen. The particle derives its name from the fact that it has no electrical charge; it is neutral. Neutrons are extremely dense.
Symbol: n^0
mass: 1.675 x 10^-27
Radioactive
When Atoms have more protons then neutrons so therefore become unstable
Isotopes
Isotopes are atoms that have the same number of protons (same) but different number of neutrons
Nucleus
The nucleus is the positively charged center of an atom and contains most of its mass. It is composed of protons, which have a positive charge, and neutrons, which have no charge.
Atomic Number/What Letter represents Atomic Number
Number of protons a given atom carries / The letter “Z” represents the atomic number
Mass Number/What Letter represents Mass Number
Number of nucleons (Number of protons + Number of Neutrons) in a nucleus
The letter “A” represents the Mass number
Photons
a quantum of electromagnetic radiation/Em radiation is a stream of light energy particles. Each of these particles are known as photon
Orbital
A region around the nucleus where electrons are highly likely going to be found. Their shapes depends on the l value
Quantum Numbers
The set of numbers used to describe the position and energy of the electron in an atom are called quantum numbers.
Principal Quantum number (n)
It describes the shells or energy levels which electrons can exist within an atom. The distance between each level and the nucleus is based on the amount of energy the electrons have./ n = the energy level around the atom, basically each electron shell = 1n, 2n, 3n
***Electrons can’t exist between energy levels(No decimals/fraction values)
Secondary Quantum Number (L)
the shapes of atomic sublevels and their orbitals.,
shapes of atomic sublevels and their orbitals.
For energy level n, the available values of L range from 0 to (n-1). For example,
n = 1, L = 0
n= 2, L = 0, 1
n= 3, L = 0, 1, 2
The values of l are more known by their letter designations.
l Letter Designations
0 s (sharp or spherical)
1 p (principal)
2 d (diffuse)
3 f (fundamental)
4 g
Magnetic Quantum Number (mL)
The orientation of the sublevel shapes in space. The values can range from -L to +L
For example, for L = 2, the ml available are: -2, -1, 0, 1, 2. Each value signifies a different orientation of the sublevel.
The quantum spin number (ms)
The quantum spin number denotes the two possible states of the electron when it is placed in an external magnetic field.
The two possible states are +½ and -½ , where they represent the spin directions of the electron.
Energy levels (shells)
the orbits around the nucleus.
Sublevels (subshells)
the different shapes and energy of electron orbitals can take on in a certain energy level. (s, p, d, f, and g)
Electron Configuration
to show how electrons can arrange themselves within an atom in the ground state (the lowest energy state) which showcases the patterns within the periodic table.
The electron configuration will give the location and the number of electrons that are found in the energy levels of a neutral atom, or an ion.
Aufbau Principle
The Aufbau principle describes that an atom can be seen as being assembled by filling it up with electrons.
It will always fill electrons beginning at the lowest available energy orbitals before filling the next higher energy orbitals.
This, in turn, provides us the sequence of electron configuration:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s
Pauli-Exclusion Principle
“No two electrons can ever have the same set of quantum numbers.”
This means that two electrons can have the same n, same l, and same ml for as long as the two electrons have different (or opposing) spin numbers.
In other words: “Each orbital (each ml orientation) can house two electrons.”
Hund’s Rule
Hund’s Rule is used when we have partially filled sublevel orbitals. Hund’s rule is stated as follows:
“Every orbital is singly occupied before it is doubly occupied.”
This means that there is a certain sequence which electrons will fill into orbitals within a sublevel.
RadioIsotope
an isotope that spontaneously decays to produce two or more smaller nuclei and radiation
How does Dalton’s model describe atoms?
Small indestructible spheres
stated:
- All matter is made of atoms.
- Atoms of the same element have the same average mass, size and unique properties.
- Atoms cannot be converted into atoms of another element through chemical reactions.
- Atoms of different elements combine in specific proportions to form compounds.
How did Thompson prove that electrons exist and that they are negative?
Experimented with cathode ray tubes which create beams of small particles. He placed charged plates on either side of the beam to determine its charge.
Observed that the beams are deflected towards positively charged plates. The beam must contain negative particles because opposite charges attract.
Since atoms are neutral, they must contain negative particles that can be removed from a positive material.
Describe Thompson’s model
Atoms are positively charged spheres with embedded negative electrons (Plum Pudding Model).
How did Rutherford prove the nucleus exists and that it is positive?
Shot positively charged alpha particles from radioactive elements at a thin piece of gold foil.
Observed that most particles went straight through. Some deflected at large angles back towards the particle gun.
Since most alpha particles went straight through, atoms are made of mainly empty space. The ones that deflected back must have hit a dense positive structure because they were repelled off.
Describe Rutherford’s model
Atoms have a small dense positive nucleus with electrons travelling around it (planetary model). He concluded that the nucleus was made of individual particles (which he called protons) equal in charge to that of a negative electron.
Describe the inferences Bohr made based on Rutherford’s model
Because of opposite charge attraction, electrons should quickly spiral into the nucleus causing a nuclear explosion.
Since this is not often observed, electrons must behave differently than macroscopic objects because they are so small.
400BC?
Democritus suggest the existence of atoms, based on intuition
no proof for 20 centuries until tools were developed
1803 - John Dalton?
“Billiard Ball” model - matter is composed of tiny particles called atoms
1) all matter is made of atoms, atoms are indivisible and in destructible
2) All atoms of an element have the same mass and properties
3) compounds are made by combing of two or more different kinds of atoms
4) a chemical reaction rearranges atoms
1904 - J.J Thomson
“Plum Pudding” model - large sphere of positive, with smaller charged particles inside it
since atoms are neutral it was believed that electrons are negative particles imbedded in a positive ball
he discovered negative charged particles
used a cathode ray tubes that was negative and positive side he shot particles through the negative to the positive, The rays did not hit the negative side and curved to the negative size
Dark ages?
The idea of atoms was frowned upon. not much progress made not really good time for science
1911 - Ernest Rutherford
“Solar system” model - small positively charged nucleus most of the mass.
based on the fact that the nucleus had a positive charge, he assumed the protons were located in the nucleus
Problems
- how does nucleus stay together?
suggestion of neutrons to keep them together
- Why does the electron not fall into the nucleus and destroy the atom
Gold foil experiment
Shot alpha particles through gold foil(Most went through) suggestion atom is mostly empty space
some alpha particles were deflected or bounced back
Light?
Matter was thought to be composed of particles, mass and specific position
300 BC: Light is particles
1600: Lights are waves
newton: light is particles
- Light was thought to be an electro magnetic wave that had mass or position
- Low frequency waves have long wave length
- high frequency waves have short wave length
Light - a electromagnetic wave of composed continuous wave lengths that form a spectrum
Photoelectric Effect
Photoelectric effect: Electron are emitted by matter that absorbs energy from shortwave electromagnetic radiation.
simple def - Shining light on metals causes electrons to emit from illuminated surfaces
- The intensity(brightness) of the light did not affect the kinetic energy of the electron ejected. The frequency of the light determined the energy of electron emitted
Max Planck
He proposed that light energy is “quantized” which meant that light was released in small discrete packets of energy, not a continuous stream
Light is provided by the light is dependent on it’s frequency
Energy gained and lost r lost or gained in whole numbers
E = nhf
E = energy in joules
n = is any integer
h = h is Planck’s constants (6.63 x 10^-34J’s)
f = frequency of the radiation
Quantum
Unit or packet of energy
smallest unit of money is a penny
Bohrs Rutherford’s model
Combination of rutherfords idea of the nucleus and bohrs founding of the electron
Accurate in it’s prediction for energy levels
However the electron energies predicted were not consistent
Cause orbits do not exist Quantum model is better
Atomic orbitals are mathematical constructs and strictly speaking are only genuine wave functions in one-electron systems such as the hydrogen atom.
Quantum mechanics
The application of quantum theory to matter
Erwin Schrodinger
Schrodinger Wave Equation allows us to describe different waves which electrons can exist within the orbital within the atom. The orbitals are differentiated from a set of 4 quantum numbers where they describe different properties within a orbital
d-orbital behavior?
Notice that on the graph of relative energy for each sublevel, 3d and 4s energy requirements are VERY close to each other. Although 4s is considered to be the higher energy level, 3d requires (a little bit) more energy than that of 4s.
This means that during a reaction, the lost electrons can come from BOTH 4s and 3d sublevels.
This is a unusual property that makes elements found in groups 3 - 12 to take on more than one ionic states.
Niels Bohr
Niels Bohr are moving stable circular orbit called energy levels (energy levels are a integer value)
electron can jump to higher energy levels by absorbing the right amount of energy
electrons from higher energy levels will return to lower energy levels by emitting excess energy in the form of light associated to the elements spectral line.
Albert einstein (1905)
Em radiation is a strem oif energy particles each of these particles are known as photons. Wave enrgy is a photon is E=hf
Electrons will only “jump” out of the surface with the right amount of energy; too high or too low
oil drop experiment?
Millikan’s original experiment, this method offered convincing proof that electric charge exists in basic natural units.
Order of Electromagnetic Waves
Radio waves - longest
Radar
Microwaves
infrared
UV
xRays
Gamma ray
cosmic rays - shortest
Longer the wavelength Shorter the Frequency
Shorter the Wave length more deadly
also *ROYGBIV
Rutherfords model
Atoms have mostly empty space
Most mass in the nucleus (Positive)
Electron float around the nucleus(no shape) and are very light
