Chemistry Unit 2 Revision

electron cloud

The region around an atomic nucleus where electrons are likely to be found.

valence

electrons involved in bonding

electorn shielding

the outer electrons are blocked from feeling the full positive charge froim the nucleus because of the inner electrons

anion

gain electrons

cation

lose electrons

energy level

a region around the nucleus of an atom where electrons are likely to be found, associated with a specific amount of energy.

sublevel

energy levels which are subdivided into sublevels, s,p,d,f

orbital

a region where the electron is found, there is one orbital for 2, three for p, etc.

electron configuration

arrangement of electrons int the energy levels, subshells, and orbitals, (electorn cloud) e.g. 1s²2s²3p^6, etc.

atomic radius

how big an atom is (distance from center of the nucleus to outermost valence electrons)

first ionization energy

energy required to move an electron

electronegativity

tendency to attract electrons to itself

group/family

a coulumn on the periodic table

period

a horizontal row in the periodic table

ground state

the lowest energy state of an atom

excited state

the energy state of an atom when it has absorbed energy and its electrons have moved to higher energy levels.

photon

a type of quantum; tiny particle of light and electromagnetic energy. (energy that travels through space in waves)

effective nuclear charge/attraction

metals/non-metals/metalloids(semi-metals)

atomic mass

measured in amu, protons + neutrons in the nucleus of an atom.

atomic number

the number of protons in an atom

Periodic trends

AR - increases down the group, and right to left (to the left) of the periodic table

EN & IE increases up the group, and to the right of the periodic table

Hund’s rule of Multipicity

within a sublevel, e’s go in separately (singly) with the same spin and pair only when needed (which checkout electron is going to)

Aufbau Principle

lowest energy orbital fills first (like a pool level represents each orbital and it must fill to reach to a higher sublevel)

Pauli Exclusion Principle

within an orbital, the two e’s must spin opposite ways (to prevent violating quantum rules)

Mendeleev

organised by atomic mass

Moosely

organised by atomic number

overtime the atomic model

became more and more complex as scientists discovered more and more parts about atoms and subatomic particles

quantum numbers

a set of unique numerical values for each electron which are depicted by the energy levels, sublevels, orbitals, and the electron spin of each electron.

only used on ground state diagrams (for now)

1st, energy level

n; s = the different levels of each electron (e.g. 2p means n = 2) etc. describes the size and energy of the orbital that the electron is in; principal quantum number; the further the distance is (down the energy levels) the more energy required for the electron to be there (nucleus is positive, electron is negative, so it is weaker attraction (magnets being close, vs far), the electron needs to work against the attraction

higher energy level - depicts size and energy, because the further the energy level the greater the distance between the nucleus and the electrons, and the greater energy required for the electrons to stay that way, and the weaker attraction (like a magnet)

2nd sublevel

the angular quantum number describes the shape of the orbital that the electron is in, which is the sublevel. l; l = the sublevel the electron is in. for s - l = 0, for p - l = 1, for d - l = 2, for f - l = 3

the different sublevels represent the different shape of the electron because of the different orbitals

3rd - orbitals

the magnetic quantum number; ml defined the orientation of the electron in an orbital.

s = 0

p = -1, 0, 1

d = -2, -1, 0, 1, 2

etc.

4th spin

ms; the direction which the electron is spinning, either up or down; +1/2 = up, -1/2 = down (rotational movement of an electron)

parts of the periodic table

1. alkali metals, 2. alkaline earth metals, 3. transition metals, 4. semi-metals (metalloids), 5. non-metals, 6. halogens, 7. noble gases 8. lanthanides, 9. actinides (8,9 are rare metals)

left is metal, right is nonmetal

when drawing excited state

move any electron to any higher level, make sure to draw circle to arrow.

energy of a photon

E = hc/wavelength = (planck’s constant) x (speed of light) / wavelength]

also equals e = hv, v being frequency bcoz c = wavelength x frequency, but that’s irrelevant as of right now

plancks constant

6.626 × 10^-34 J x s

speed of light in a vacuum

2.998 × 10^8 m/s

if wavelength is in nanometers

then to convert to joules, must add 10^-9 to the wavelength value, e.g. 461 nm would be 461 × 10^-9

to calculate wavelength

1/wavelength = -0.01097nm^-1[1/n² high - 1/n² low]

high = the higher energy level (high)

low = the lower energy level (low)

for a photon to be visible, and the wavelength determines, atomic spectrum

wavelength determines colour of the light, and the photon to be visible, the photon must have a wavelength for 400 -700nm. the observed wavelengths create a pattern of light called the atomic spectrum

relaxation

the process by which an excited electron returns to a lower energy level, emitting a photon in the process (may or may not be visible)

photoelectric effect

electron absorbs sm energy that it moves past the excited state and becomes ionized as a cation.

1st; speed = distance/time, c is speed of light, so it is also c = distance/time.wavelength is the distance while the frequency (v) is per time (nm) (1/sec) = nm/sec. E = mc², so MC = h/wavelength

s orbital

holds up to two electrons

p orbital

dumbbell shape; each can hold up to 2 electrons, but all shape of three orbitals hod up to 6 electrons, x = horizontal, y = vertical, z = diagonal

this is the complete one which holds 6

d orbital

dumbbell shape with a ring; contains five orbitals, each can hold up to 2 electrons, allowing for a total of 10 electrons.

f orbital

complex shapes; contains seven orbitals, each can hold up to 2 electrons, allowing for a total of 14 electrons.

history of an atom

Solid Sphere, Plum Pudding, Saturnian, Nuclear, planetary (modern day model)

Atomos

Democritus; 400 BC; atoms are the building blocks of matter, the shape of an atom explain elements behavior.

Solid Sphere

John Dalton - 1803, solid, uniform, indivisble mass with no internal structure

Plum Pudding

J.J. Thompson - 1904, with a positive charge in the middle and electrons negative charge surrounding it/ negative electrons embedded in a sea of positive charges

Saturnian

Hantaro Nagaoka - 1904, positive charge in the middle and negative charges (electrons) around the positive charge like Saturn rings,

Nuclear

Ernest Rutherford - 1911 where the electrons are surrounding the central nucleus (positive charge)

Model Day (Planetary)

Niels Bohr (electrons on fixed circular orbits or shells surrounding on a central nucleus) - 1913 2,8,2

Quantum

1926 - Schrodinger, electrons are in clouds, surrounding the nucleus, and this cloud is less dense.

left to right a period the attraction

increases because the number of protons increase which creates a higher positive charge for the nucleus, making attractive forces higher, pulling the electrons closer to the nucleus.

moving up a group/family, the attraction

increase because the energy levels decease meaning that there is less electron shielding, attraction forces increase because the nucleus is closer to the electrons

the greater the distance

the weaker the attraction force, the atomic radius is larger, but the lower ionization energy and electronegativity (because the attraction of the electron to itself is lower); electron shielding increased which means that the valence electrons are shielded from the nuclear charge due to inner electrons.