Unit 1 Chemistry Structure and Properties of Matter

Democritus

Theorized that particles get to a size where they can no longer be divided. He called these particles atoms, Which means indivisible (This is where we got the name the atom)

John Dalton

Atoms are small indestructible spheres

“Billard ball” model - an atom is a solid, uniform, sphere

J.J Thomson

Used the cathode ray tube experiment

• shot a high voltage through two electrodes at one end of a tube which allowed of a beam of particles to flow from the cathode ray ( - charged) to the anode ( + charged)

• To test the properties of the particles, He placed two oppositely charged plates around the cathode ray and the cathode ray was deflected

• This led him to the theory that a atom consists of negative charges

• “Plum pudding” mode - The atom is a positively charged mass with discrete negative particles throughout

Ernest Rutherford

Created the gold foil experiment

• Positively charged alpha particles were shot at a very thin sheet of gold foil

• Most particles went straight through, but some particles were deflected, and a few bounced back

• The “nuclear” model - atoms are mostly empty space with a small dense, positively charged “nucleus”, electrons move in the empty space that makes up the rest of the atom

What are some of the limitations of his theory?

• Nucleus composed of only positive charges should break apart due to repulsion force

• Could not explain the total mass of an atom

Neils Bohr

Spectral Line experiment - High voltage connected to terminals of a tube containing hydrogen gas

• When electricity hits an electron, it gives it some energy moving it further from the nucleus to an “excited state” the electron then spontaneously drops down to its normal energy level ( The ground state )

• When an electron returns to its ground state, it gives off an amount of energy which can be sometimes seen as a color of visible light

• When emitted light was through a prisim, only fourth distinct colors of light (line spectrum) were seen

• Bohrs atomic model - Postivley charged nucleus with electrons travelling around it in specific energy levels

• “Quantized energy” - Because energy is restricted to only certain discrete quantities

What are some limitations with Neils Bohrs theory

• Theory could only explain line spectrum of hydrogen but not accurately predict the spectral lines of multi electron atoms

What is the quantum mechanical model of the atom?

• Describes the three dimensional position of the electron in a probabilistic manner according to a mathematical function called a wave function

What did Louis Debrogile come up with?

Matter consists of both a particle nature and a wave

What did Erwin Schrodinger come up with?

• Expressed the behavior of electrons as matter waves

• Came up with the quantum mechanical model

What did Schrodingers wave function show?

1. Electrons exist in defined energy levels. The energy of the level increases with the number. Numbers are always whole numbers

2. Electrons done exists in defined orbits, but rather in clouds of probability, as seen in the figure above, called orbitals

State The heinsbrurg uncertainty principle

• States that both the position and the velocity of an object cannot be measured exactly, at the same time, even in theory

What is an orbital

An orbital (rather an orbit) is the term used to describe the region of space where electrons may be found. Sometimes called “electron clouds”

What is the principal quantum number (n)

• Denotes the energy level or shell, of an atomic orbital and its relative size

• A higher value indicates a higher energy level that is larger in size and further away from the nucleus

What is the secondary quantum number (l)

• Refers to the orbital shape or subshells, within each principal energy level

• The numbers of L are usually replaced by letters to avoid confusion with n (s,p,d,f)

What is the magnetic quantum number (m1)

• Describes the orientation of the orbital in the space around the nucleus

• It has whole number values between - L and + L

• If L = 1 then m1 can be -1,0,+1

What is the spin quantum number (Ms)

• The spin quantum number describes the spin of an electron

• Two values only

What is Paulis exclusion principle?

• No two electrons can have the same quantum numbers

• No two electrons can have the same magnetic spin

What is Hund’s rule

• One electron occupies each of several orbits at the same energy before a second electron can occupy the same orbital

What is Aufbau Principle

• Each electron is added to the lowest energy orbital available in an atom or ion

What is electronegativity

The relative ability of an atom to attract shared electrons in a chemical bond

Tell me the ranges of electronegativities

EN 1.7-3.3 = Mostly Ionic

EN 0.4 - 1.7 = Polar Covalent

EN between 0.0 and 0.4 = Mostly covalent

What is metallic bonding?

• Chemists use the electron sea model to describe metallic bonding. The model proposes that the valence electrons of metal atoms move freely among the ions, forming a sea of de-localized electrons that hold the metal ions rigidly in place

What are some properties of metallic bonds

• The stronger the bonding forces, the higher the melting / boiling points of pure metals

• Metals are good conductors because their electrons are free to move from one atom to the next

• When struck metal ions can slide by another while the electrons still surround them making them malleable

What are alloys?

• Alloys are solid mixtures of two or more metals

What is ionic bonding

Essentially involves one atom losing one or more electrons and another atom gaining the electrons

Properties of ionic compounds

• They have high melting and boiling points due to very strong attraction between ions

• Ionic compounds are soluble in water when the attractive forces between the ions and the water molecules are strong than the forces between the ions themselves

• Solids do not conduct because ions have no room to move

• Compounds conduct when dissolved so ions can move around

What is covalent bonding

• In most cases the sharing of electrons allows the atoms to obtain a noble gas configuration

• Split into two types

  • Polar covalent : Atoms do not share electrons equally

  • Non polar covalent : Atoms share electrons almost equally

Give me the properties of molecular solids

• Form a crystal lattice structure held by inter molecular forces that are weaker than ionic or covalent bonds

• Relatively low melting point due to weak inter-molecular forces

• Non conductors because individual particles are neutral molecules

Explain what covalent network solids are

Substances that consist of atoms covalently bonded in a continuous two or three dimensional array

Properties of covalent network solids

• Among the hardest materials on earth

• Much higher melting point than ionic due to interlocking structure and many covalent bonds

• Non conductors because electrons are held within atoms or in covalent bonds

What is a co ordinate covalent bond?

• One atom contributes both electrons to the shared pair

• Bonds behave the same way as other covalent bonds and therefore are not indicated in the Lewis structure

Intramolecular forces

• A force that holds atoms or ions together, including ionic, covalent bonds, and metallic bonds

Breaking and forming intramolecular forces are…

Chemical changes

Intermolecular Forces

• Weak force of attraction that exists between two or more molecules

What are the three types of intermolecular forces

• Dipole to Dipole

• Hydrogen Bonding

• London Dispersion Forces

Explain what London Dispersion forces are

• The London force is due to the simultaneous attraction of the electrons of one molecule by the nuclei in the surrounding molecules

• Constant movement of electrons in a molecules causes a temporary dipole that includes another instantaneous temporary dipole in an adjacent molecule; The process “disperses” through the substance, creating fleeting dipoles that attract each other

London Dispersian forces are the primary force of attraction between

non polar molecules such as hydrocarbons

When is the London dispersion force strong

• When the molecule is bigger because they have more electrons which results in a greater shift in electoral charge and stronger temporary changes

• When there is a larger surface area

Explain what Dipole to dipole forces are

• Strong force of attraction between two or more polar molecules

• The partial negatives and positives will attract each other

• These are stronger than the LDF and permanent

Explain what hydrogen Bonding is

• A very strong form of dipole to dipole attraction

• Occurs when a molecules has a hydrogen atom bonded to a highly electronegative atom, often an O,F,N atom

  • The strongly electronegative atom pulls the shared electron pair away from the hydrogen atom

  • The hydrogen atom will strongly attract to another lone pair of electrons on another O or N of a nearby molecule

Explain how surface tension works

• Molecules within a liquid are attracted by molecules on all sides, but molecules right at the surface are only attracted downward and sideways. Therefore, liquids tend to stay together. Water has stronger inter molecular forces, and therefore has a very high surface tension

Give me 4 reasons on why hydrogen Bonding is amazing

1. The surface tension

2. Relative high boiling point of water

3. Giving solid water (ice) a lower density than liquid water because of its bent shape which gives a great deal of space between each molecule

4. For waters ability to act as a solvent - Polar substances easily dissolve in water