Atomic Structure Last Unit of Grade 12 Chem!!!!

Electromagnetic Spectrum An element’s chemical behavior is related to the arrangement of the electrons in its atoms. Scientists have discovered there is a link to atomic structure and the electromagnetic spectrum.

The electromagnetic spectrum consists of electromagnetic radiation, which is the release and transmission of energy in the form of electromagnetic waves.

Electromagnetic Radiation Is a form of energy that exhibits wave-like behavior as it travels through space

Examples: Visible light from the sun Microwaves that heat or cook food X-rays to examine bones and teeth TV and radio waves

Wavelength (λ, lambda) The shortest distance between equivalent points on a continuous wave

Frequency (ν, nu) Is the number of waves that pass a given point per second Measured in hertz (Hz) or waves per second

Amplitude The wave’s height from the origin to a crest or trough

Although the speed of all electromagnetic waves is the same (3 Although the speed of all electromagnetic waves is the same (3.00 x 108 m/s [speed of light included]), waves may have different wavelengths and frequencies

Wavelength and frequency are inversely related Wavelength and frequency are inversely related. This means as wavelength increases, frequency decreases and vice versa.

Calculations C=λν Where c = the speed of light (3.00 x 108 m/s) λ = wavelength (in meters) ν = frequency (in Hz)

Example What is the wavelength of a microwave having a frequency of 3.44 x 109 Hz? Use C=λν to solve. Answer: λ = 8.72 x 10-2 m

Energy Max Plank found that the amount of energy absorbed or emitted by a body is proportional to the frequency of the radiation. E = hv Where: E = energy in Joules (J) v = frequency (Hz) h = 6.63 x 10-34 Js or J/Hz

Example Calculate the energy of a wavelength with a frequency of 5.00 x 1015 Hz. Answer: 3.32 x 10-18 J

Electromagnetic (EM) Spectrum Includes all forms of electromagnetic radiation, with the only differences in the types of radiation being their frequencies and wavelengths

Lower frequency vs Higher frequency

Visible light is the portion of the EM spectrum that is detected by the human eye with its colors ranging from red though violet (ROYGBIV)

The energy of the various parts of the EM spectrum is directly related to the frequency of the wave. ie. if a wave has a high frequency, then it will contain a higher amount of energy.

Line Spectra Emission spectra can exist as a continuous spectrum (a rainbow) or a line spectrum.

A line spectrum (atomic emission spectrum) consists of distinct bright lines appearing on a dark background that occur in different parts of the visible spectrum.

It is this distinguishing feature of gaseous atoms that provide us with a unique “fingerprint” for each element.

Each element has its own unique line spectrum as each of these elements contains differing amounts of electrons and different energy levels.

These bright lines indicated that only certain energies are possible within the atom. The brightness of spectral lines depends on how many photons of the same wavelength are emitted.

Photons Particles of electromagnetic radiation with no mass, that carry a quantum of energy (light energy)

Quantum The minimum amount of energy that can be gained or lost by an atom

Spectral lines are produced by an atom in the excited states Spectral lines are produced by an atom in the excited states. First, electrons absorb energy which causes them to rise to a higher energy level.

When the electron falls back down to a lower energy level, it gives off a color of light. Since there are many electron transitions possible between energy levels, there are many spectral lines produced by an atom in the excited state.

Example: When the electron falls from the following energy levels, the corresponding colors are emitted:

Hydrogen: Third to second  red (this is the shortest distance, thus lowest frequency of visible light) Fourth to second  green Fifth to second  indigo Sixth to second  violet

Where have you seen line spectra? Neon lights Fireworks Aurora borealis (Northern Lights)

Electron Configuration The arrangement of electrons in an atom There are three principles (rules) that define how electrons can be arranged in an atom’s orbital.

The Aufbau Principle States that each electron occupies the lowest energy orbital available. eg) 1s orbital is filled before the 2s orbital

All orbitals in the related sublevel have the same energy eg) all three 2p orbitals are of equal energy

The energy levels within a principal energy level have different energies. eg) the three 2p orbitals are of higher energy than the 2s orbitals

Orbitals related to energy sublevels within one principal energy level can overlap orbitals related to energy sublevels within another principal level. eg) the orbital related to the atom’s 4s sublevel has a lower energy than the five orbitals related to the 3d sublevel

The Pauli exclusion principle States that a maximum of two electrons may occupy a single atomic orbital, but only if the electrons have opposite spins  The atomic orbital containing paired electrons with opposite spins is written

Hund’s Rule States that single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbital

The sequence in which 6 electrons occupy the three P orbitals is:

Zigzag Rule

Atomic Radius The atomic radius is simply the distance from the nucleus to the outermost electrons. Since the edge of the outermost electrons can never be known precisely, the atomic radius is usually defined as ½ the distance between the nuclei of atoms of the same element.

1) Periodic Trends (Rows) The atomic radii generally decreases as you move across a period (row). Since each additional electron is added to the same principal energy level, the additional electrons are not shielded from the increasingly positive nucleus. (number of protons) The increased nuclear charge pulls the valence electrons closer to the center of the atomic nucleus

2) Group trends (columns) The atomic radii generally increases as you move down a column. As you move down a column the outermost orbital increases in size shielding the valence electrons from the center of the atom. These factors overpower the increased pull of the more positive nucleus on the valence electrons causing the atomic radius to increase.

1) Periodic trends The size of positive ions decreases as you move across a period and the size of negative ions increases as you move across a period.

2) Group trends The radii of both positive and negative ions increase as you move down a group.

Note: the size of an atom decreases when being converted to a positive ion because it loses an electron and therefore there is less electron repulsion. The size of an atom always increases where being converted to a negative ion because there is an increase in repulsion between electrons.

Ionization energy The energy required to remove an electron from a gaseous atom These values indicate how strongly an atom’s nucleus holds onto its valence electrons. High ionization  strong hold on its electrons Low ionization  loses its outer electrons easily

1) Periodic trends Ionization energy increases as you go left to right across a period. This is because of the increased nuclear charge (as the number of protons increases) which holds the electrons more strongly, making it more difficult to remove the electron.

2) Group trends Ionization energy decreases as you move down a group. The increased atomic size pushes the valence electrons further from the nucleus, so it takes less energy to remove the electron because the strength of attraction is decreased.

Successive ionization energies After removing the first electron from an atom, it is possible to remove additional electrons as shown in the table below of period two elements. Notice how the energy required for each successive ionization always increases, but not always smoothly.

For each element, there is a dramatic energy jump related to that element’s valence electrons. Electrons hold onto their inner core electrons much more strongly than their valence electrons. For example, oxygen has six valence electrons so the ionization energy is much higher than the first energy.

Electronegativity Indicates the relative ability of an element’s atoms to attract electrons in a chemical bond

Periodic trends Electronegativity increases as you move from left to right across a period. This is because there is greater electron-affinity power of the nucleus with the increased nuclear charge as the number of protons increase from left to right.

Group trends Electronegativity decreases as you go down a group. This is due to the shielding where electrons in lower energy levels shield the positive charge of the nucleus from valence electrons resulting in those outer electrons not being as tightly bound to the atom.

Electronegativity differences By taking the difference in electronegativities of two given atoms in a polyatomic compound, the bond character can be predicted using the table below.

Electronegativity differneces Electronegativity Difference 0.0 – 0.4 0.4 – 1.0 1.0 – 2.0 ≥ 2.0 Bond Type Non-polar covalent Moderately polar covalent Very polar covalent Ionic

Example: What type of bond would occur in a molecule of lithium fluoride? Ionic

From the electronegativity table: Li has a value of 0.98 F has a value of 3.98 Difference: 3.00 According to the table above the bond between Li and F is ionic