Grade 11 Chem Unit

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Empirical Knowledge

Empirical Knowledge

Comes from observations and experementations.

Theoretical Knowledge

Based on knowledge created to explain observations

Theory

Comprehensive set of ideas that attempt to explain a law or related observations.

Model

Mental or physical representation of a theoretical concept

History of the Atom

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Timeline and explanation

1. Empedocles: everything’s made up of 4 elements(water, earth, air, and fire)

2. Democritus: geometric shapes

3. Dalton: billard ball

4. J.J. Thomson: the muffin(cathode ray experiment)

5. Rutherford: intorduced nucleus with protons(gold foil experiment)

6. Bohr: introduced orbitals/ diff energy levels

7. Chadwick: introduced neutrons(where its weight is added)

History of the Periodic Table

1. 330 B.C → 4 elements

2. 1800 → 31 Elements

3. 1865 → 61 Elements

John Alexander Newlands (1864): Law of Octaves made

Dimitri Mendelev(1869): Father of Periodic Table, created the Periodic Law

Henry Mosely(1911): patterns seen when [ut in orderof atomic number + grouped them into diff families and periods

Law of Octaves

every eighth element has similar properties as the 1st one

The Periodic Law

When arranged by by atomic number, the properties of elements repeat at regular intervals.

Lanthanids

• Rare earth metals

• elements 57-70

Actinides

• Radioactive

• Elements 89-102

• Transuranic Elements: sythesis elements (93+)

Isotope

atoms of the smae element that have diff masses due to diff #of neutrons

eg. oxygen-17

Format→X-#

Average Atomic Mass Formula

Atomic Mass = Atomic mass * abundance % + (other Isotopes)

Mass measured in u or amu

% Abundance represented as a decimal

Isotope Types

1. Stable: don’t naturally decay but can exist in natural materiels in diff proportions

2. Unstable: continuously and spontaniously break down/decay in lower atomic weight isotopes.

   1. releases readiation

   2. neutrons leave

Half-life

the time it takes for ½ of the parent mass to decay into the daughter mass

Trends of the Periodic Table

1. Atomic Radii

2. Ionic Radii

3. Electronegativity

4. Ionization Energy

5. Electron Affinity

Atomic Radii

half the distace form the nucleus of 2 identical bonded atoms(mesured in picometers)

Trend:

• Increase Down a group(less attraction cuz of more e -)

• Decrease across a period(more attraction cuz more p+)

Ionic Radii

Cation: have smaller ionic radius than corresponding atom(p+ > e -), less sheilding of e -, e - pulled closer

Anion:have larger ionic radius than corresponding atoms(e - > p+), greater electron-electron repulsion

• Ion size increases down a group

Elecronegativity

the ability to attract electrons when bonded

Trend:

• Increase across each period:since the bonding pair can get closer to the nucleus and be more attracted

• Decrease down a group: since atomic radius is larger, the atoms attract bonding pairs less strongly

non-metals have a high electronegativity

metals have low electronegativity

ionization Energy

the energy required to remove an electron from a neutral atom of an element(mesured in kilojoules/mol)

Li + energy →Li+ + e -

Trend:

• Increase across each period: since atoms are getting smalling, electrons are closer to the nucleus

• Decrease down a group: atoms are getting bigger, e - are father from the nucleus, outer electrons become increasingly more sheilded from the nuleus by inner electrons(less energy needed for e - to be taken away)

Metals have Low IE

Non- Metals have high IE

Noble gases have VERY High EI

Electron Affinity

the energy that’s released when a neutral atom aquires an election

eg. Cl + e - →Cl 1- + energy

if an atom is forced to take an electron, ot much energy is realesed due to the ammount of energy used to force the e - in

eg.Na + e - + energy→Na+

Trend:

• Increase across each period: Halogens more prone to gaining electrons

• Decrease down a group: more prone to losing electrons