unit 1: matter, chemical trends, and chemical bonding

Democritus

-first proposed indivisible particles called atoms

Dalton’s Model

• Called the “billiard ball” model

• Atom was a solid sphere

• Atoms of same element are identical

Thomson’s Model

• Called the “plum pudding model”

• Discovered electrons

• Atom was a positively charged sphere, with negatively charged particles

rutherfords model

• Discovered nucleus

• Used Gold foil experiment

  • piece of gold foil was hit with positive alpha particles

  • Most alpha particles went straight through

  • Showed that gold atoms were mostly empty space

• Planet-like electrons orbit a positively charged nucleus

chadwick

discovered the neutron

Bohr’s Model

• Electrons only have a specific amount of energy, 

• Organized in energy levels called shells

• Electrons gain/lose energy to move between shells

Schrodinger

• Showed that electrons move in a region of space, often represented as a cloud

explaining flame tests

When an atom is subjected to heat or electricity, the electrons in the atom become excited

The electrons absorb energy and jump from the ground state to the excited state.  This is called a quantum jump.

What goes up must come down.  Eventually the electron in the excited state will fall back to the ground state and release energy in the form of coloured light.

quantum jump

The electrons absorb energy and jump from the ground state to the excited state.  This is called a quantum jump.

how different colours of light are created

The colour of light emitted depends on how big the jump is and thus the amount of energy released.

 

remember for explanation:

• large/small jump

• high/low energy

• smaller/larger wavelength

line spectrum

• a series of coloured lines separated by bands of blackness

• when the coloured light is viewed through a spectroscope 

• unique to every element

• like a fingerprint that aids in identifying the element

isoelectronic

• Noble gases have 8 valence electrons and therefore a very stable shell

• Helium is the exception as it is stable with only 2 valence electrons (first shell is full with 2)

• Some atoms acquire stable octets by gaining or losing electrons.  

Once they acquire a stable octet (usually 8 in an outer shell), they are said to be isoelectronic with the noble gas that has the same total number of electrons.

protons

• positively charged particles in the nucleus

• Determine the identity of the atom

electrons

• Negatively charged particles

• Orbit the nucleus

• Transferred in chemical reactions

isotopes

Atoms of the same element that have different masses due to a different number  of neutrons

• same number of protons and electrons means similar chemical/physical properties

radioisotope

• An unstable isotope of an element

• Undergoes radioactive decay into more stable nuclei

why is atomic mass a weighted average

The existence of isotopes can explain why the atomic mass on the periodic table is an average atomic mass.

This mass is a weighted average.

chemical bonds

The forces that attract atoms to each other in compounds

electronegativity definition

A measure of an element’s ability to attract electrons in a chemical bond

electronegativity

• If the electronegativity of one of the two atoms in the bond is greater than the electronegativity of the other atom, the electrons will be more strongly attracted to the first atom

• Electrons spend more time around atoms with higher electronegativity

how to determine type of bond 

• Compare electronegativity values of two elements in a chemical bond

• This determines whether they will equally share electrons (COVALENT), share electrons unequally (POLAR COVALENT), or share so unequally that they transfer electrons (IONIC).

most polar to least polar

polar

• uneven distribution of charge 

has a partially positive (∂⁺) and partially negative (∂^-) end (dipole)

* the atom with higher delta EN is partially negative, lower delta EN is partially positive

keeping molecules together

Opposite ends of partially charged molecules attract creating intermolecular forces

The higher the ∆EN, the higher the melting/boiling points

formation of ionic compounds

• Ionic compounds transfer electrons to produce positive ions and negative ions

• The opposite charges attract 

• Charged ions conduct electricity when dissolved in water

• The attraction is strong and therefore ionic compounds:

  • have HIGH melting point 

  • are solids at room temperature

lewis structure for molecular compounds

1. Determine, from the chemical formula, the number of atoms of each type of element in the compound.

2. Use the periodic table to determine the number of valence electrons for each atom.  Add these up for the total number of valence electrons in the compound.

3. The element that requires the most bonds is the central atom.  Arrange the other elements around the outside.

4. Place two dots between elements (bonding pairs) 

5. Place remaining valence electrons around the atoms in pairs (lone pairs) to complete their octets, do the central atom last

6. If there are not enough electrons for all atoms to have 8, create double or triple bonds by sharing additional electron pairs

the modern periodic table

• Modified Mendeleev’s table

• Organizes elements according to atomic number

periodic law

Chemical and physical properties of elements repeat in a regular pattern when arranged by increasing atomic number

periods, families (groups), valence electrons

Periods: horizontal rows

Groups (aka families): vertical columns, same number of valence electrons gives similar properties

Valence electrons: electrons in the outermost energy level (orbit), involved in chemical reactions

when examining periodic trends…

When examining periodic trends, always look at:

• # energy levels (orbits)

• # protons

As we go down a group: # energy levels increase

As we go across a period (left to right): # protons increases

atomic radius

the size of the atom

Ex explaination: Which would be larger - Be or Mg?

• Be has 2 energy levels

• Mg has 3 energy levels

• Mg is larger

OR

Eg. Which would have the smallest radius: Mg or Si?

• Both have 3 energy levels

• Si has more protons to attract the electrons closer

• Si is smaller

atomic radius in general

Down a group:

AR increases (more energy levels)

Across a period: 

AR decreases (more protons, same energy levels) 

ionization equations

• Atoms will lose or gain electrons (react) so that they are isoelectronic with the nearest noble gas

• This makes the atom stable as its orbits are full

Ionization Energy

The amount of energy required to remove an electron from the outermost energy level of an atom or ion (in gaseous state)

More loosely held electrons are more Easily Removed, Lower I.E.

eg: Eg. Which atom has the smallest ionization energy: Li or Rb

• Less energy to remove outermost electron from Rb

• Negative e^- farther from positive nucleus in energy level 5

• Li attracts electrons more tightly because they are closer to positive pull of the nucleus in shell 2

eg: Na or Al?

• Na

• Same number of energy levels

• Al has more protons to attract the electrons, electrons more difficult to remove, requires more energy

ionization energy in general

Down a group: 

IE decreases (more energy levels)

Across a period: 

IE increases (more protons, same energy levels)

electron affinity

• How much more stable an atom is after gaining an electron

• Large negative numbers considered high EA

• Negative sign means energy is released (hurray!)

ex: does F or O jave larger electron affinity?

• Same number of energy levels

• F has more protons to attract same number of energy levels

• F will hold electrons more tightly, better attract new electrons

• F will have larger electron affinity

electron affinity

Down a group: decreases (more energy levels)

Across a period: increases (more protons)

electronegativity

• Ability of an atom to attract electrons in a bond

electronegativity in general

Down a group: EA and EN decrease (more energy levels)

Across a period: EA and EN increase (more protons)