Unit 1: Structure And Properties Of Matter Flashcards
Joseph Thompson
- experimented with cathode ray tubes
- creates a beam of small particles that make a flash of light when aimed at a screen
- he placed oppositely charged plates at either end to determine the charge
FOUND: the beam of particles was deflected towards the + plate
The beam must contain negative particles that can be removed from positive material
What did john dalton find about atoms?
Reacted atmospheric gases
All matter is made of atoms
Atoms of the same element have the same average mass, size and unique properties
Atoms cannot be converted chemically into other elements
Atoms form in whole number ratios
Ernest Rutherford 1911
Shot positively charged alpha particles from radioactive elements at a thin piece of gold
FOUND: most + particles went through but other deflected back causing him to believe they hit a small positive centre called the nucleus filled with protons
He was wrong because like charges repel and the nucleus could not be solid
Niels Bohr
Mathematically determined where electrons were by Einstein’s light theory and max plank’s specific energy amount theory
Atomic emission spectra
Believed electrons are only found in specific allowable energy levels that represent fixed circular paths around the nucleus
In orbit, electrons do not gain or lose energy
An electron changes orbits but emitting or absorbing energy
Ernest Rutherford 1920
Calculated the average mass of protons it would take to balance the negative charge of electrons
Found that atoms weighed significantly more than his calculations
Found there must be another particle that caused the extra weight but must be neutral in charge do that they don’t upset the charge balance in the atom
Werner Heisenberg
Proposed the uncertainty principle: it is impossible to measure the position and momentum of an electron at the same time
The more accurately one property is measured, the less accurate the other one is measured
Proved Bohr’s model wrong because he assigned fixed paths and definite energy levels
Louis De Broglie
Believed all matter behaved like waves
Developed an equation to determine the wave length of any object like an atom
An electron moving in the first energy level has a larger wave length than the radius of the atom
Meaning the electron doesn’t wave around in a circle, it’s actually outside of the orbit that Bohr said it should be in
Erwin schrodinger
Quantum mechanical model: Incorporates the wave like characteristics of electrons
Applied wave theory to Bohr’s model to determine the location of the electrons using Bohr’s energy levels
Created an equation which allowed him to predict the location of an electron using statical probability
Created the region of probability orbitals where the electron would most likely be 95% of the time
These are not the electron paths
Quantum numbers
Orbitals have a variety of sizes and shapes depending on the amount of energy the electron has
Principle quantum number (n)
Orbital size & energy level
Higher the n value, the larger the orbital and the energy the electrons have
Periodic table periods are set up by n value
Orbital shape quantum number l
Each value of l represents an energy sub level and describes a specific orbital shape L=0 s-shaped L= 1 p-shaped L= 2 d-shaped L= 3 f-shaped
Magnetic quantum number ml
Describes the orientation or plane of the orbital shape
Quantum spin number ms
Wolfgang Pauli said electrons can only occupy an orbital with opposite spins to balance their repulsive forces
Can only equal +1/2 or -1/2
Atomic orbitals and their relative energy order
1s _ 2s _ 2p _ _ _ 3s _ 3p _ _ _ 4s _ 3d _ _ _ _ _ 4p _ _ _
Pauli exclusion principle
No two electrons in an atom can have the same set of quantum numbers
Aufbau principle
Building up process
Each set if quantum numbers is the same ad the last element in the order of the periodic table with one last electron tacked on in the lowers possible energy level
Hund’s rule
Each orbital with the same amount of energy receives one electron before pairing
When electrons are added singly to orbitals they must have the same spin of +1/2
Condensed electron configurations
Take ground state electron configuration
Take previous noble gas in square bracket and add extra electrons needed
Excited state
When electrons are given energy they move into orbits of higher energy
Even though previous orbit is not filled, it jumps to a higher energy level
Ex. Copper is usually in an excited state
[Ar] 4s1 3d10
Should be 4s2 3d9
Copper is more stable with a full 3D shell and therefore becomes excited to increase stability
Intramolecular vs. intermolecular forces
InTRAmolecular: electrostatic attractions inside of compounds to hold atoms together in molecules
InTERmolecular: electrostatic attractions between 2 or more molecules of ions
Ionic bonds
Metal loses electron which requires an input of energy = less stable
Non metal gains electron and releases energy = regains stability
Ionic bonds form in three dimensions known as crystal lattices which causes the ions to be held into place which causes a loss in kinetic energy and becomes more stable
Covalent bonds
Atoms are brought close enough together to overlap 2 partially filled orbitals which creates a region where the electrons would be 95% of the time
The 2 nuclei and 2 electrons move to a position where their repulsive forces balance called optimal particle separation this restricts movement and slows down electrons causing a release in bond energy and as a result the atoms become more stable
Metallic bonds
Metals exist in pure solid form so they make bonds between themselves
Metal atoms do not gain electrons nor do they have enough valence to share to make covalent bonds
Arnold summerfeld proposed the free electron model: cations are packed into a crystal lattice held together by delocalized valence electrons which causes them to slow down and release energy
Metallic bond strength increases the smaller the metal atom and the more free electrons they can produce
Ion dipole forces
Attraction between a full charge ion and the opposite slight charge of a polar molecule
Hydrogen bond
Occurs between 2 polar molecules where o-h n-h or f-h group on one molecule attracts the lone pair of an o, n or f on a neighbouring molecule
Strong slight charge + strong slight charge
Dipole dipole
Two polar molecules where opposite slight charges attract
Ion induced dipole
An ion brought close to a non polar molecule causes electrons to displace and the non polar molecule becomes temporarily polar (iron in blood)
Dipole induced dipole
A polar molecule is brought close to a non polar molecule and causes electron to displace and the non polar compound becomes temp. Polar
LDF
Shared electrons are constantly vibrating which temporarily causes a charge imbalance
This temporary dipole can induce surrounding molecules
LDF forces increase with size and surface area
Valence shell electron pair repulsion theory (VsEPR)
Each pair of electrons repels each other so they are as far apart as possible
Linear arrangement
BA: 180
Molecular shape: linear
Triganol planar arrangement
3 electron groups
BA: 120
Shape: Triganol planar or v shaped
Tetrahedral arrangement
4 electron group
BA: 109.5
Shape: tetrahedral, Triganol pyramidal, v shaped
Triganol bipyramidal arrangement
5 electron groups
BA: 90, 120
Shape: Triganol bipyramidal, seesaw, T shaped, linear
Octahedral arrangement
6 electron groups
BA: 90, 90
Shape: octahedral, square pyramidal, square planar
Symmetrical and different arms
Polar
Asymmetrical and arms the same
Polar
Symmetrical and same arms
Non polar
Atomic solids
Made of noble gases London dispersion Extremely low melting point Extremely soft No charges so can't conduct Low solubility in water due to dipole induced dipole
Non polar molecular solids
Stacked non polar molecules
LDF
Low to medium melting point depending on size
Soft
No charges so not conductive
Low solubility due to dipole induced dipole
Ionic solids
Ionic bonds High melting point Hard Can't conduct as a solid charges can't move Ion dipole some dissolve some don't
Network solid
Atoms or compounds linked by covalent bonds Covalent bonds High melting point Hard repeated geometric pattern No charge so can't conduct Dipole induced dipole low solubility
Polar molecular solid
Stacked polar molecules so slight opposite charges attract
Dipole dipole or HB depending on compound
Medium Melting point
Medium hardness
Can’t conduct stuck in crystal lattice
Medium to medium high solubility
Metallic solids
Free electron model
Metallic bond
Medium to high melting point
Medium high hardness
Electrons can move medium to high conductivity
Ion dipole with water but doesn’t happen (forks)
