Chapter 1: Atomic Structure

Question 1

define atom

Answer 1

the smallest QUANTITY of matter that still retains the properties of matter. the most basic unit of an element.

(an element is a substance that cannot be broken down into 2 or more substances by any means)

Question 2

define sub-atomic particles

Answer 2

when an atom is divided smaller, it produces subatomic particles. its nature, arrangement, and number determine the properties of an atom and hence the matter.

protons, neutrons, electrons

Question 3

explain physical properties of subatomic particles

Answer 3

protons are positively (+1) charged particles found in the nucleus

neutrons are electronically (0) charged particles found in the nucleus. slightly larger than protons.

electrons are negatively charged (-1) particles distributed around the nucleus

electrons are significantly smaller in mass compared to protons (by 2000x), while protons and neutrons are similar in mass

Question 4

explain atomic number & mass number

Answer 4

atomic number (Z): no. of protons in the nucleus

(atoms are neutral hence the atomic number is also the no. of electrons distributed)

mass number (A): total number of protons and neutrons, collectively known as nucleons

Question 5

explain isotopes

Answer 5

atoms that have the same atomic number but different mass numbers

(most elements have two or more isotopes)

isotopes of the same element usually have similar chemical properties (same types of compounds, similar reactivities)

Question 6

explain atomic mass

Answer 6

mass of an atom in atom units (amu)

(1 amu is 1.661 x 1*10^-24 g)

relative atomic mass is the average mass of the naturally occurring mixture of isotopes

Question 7

explain groupings of elements in the periodic table

Answer 7

periods- horizontal rows, increases in atomic number

groups/family- vertical rows

group 1 (alkali metals), group 2 (alkaline earth metals), group 17 (halogens/diatomic), group 18 (noble gases)

Question 8

explain energy levels

Answer 8

arranged into main levels/shells, determined by the distance (n) of the electron to the nucleus

each energy level represents an allowed amount of energy

(electrons are quantised. they only have definite amounts of energy. they are restricted to certain energy levels.)

each shell is comprised of one or more energy sub levels/subshells: s, p, d, f

Question 9

explain energy level states

Answer 9

an atom is in ground state when the electron occupies the lowest energy level available.

when an atom absorbs a discrete amount of energy (n where it cannot be a non-natural number, only positive whole), it moves to a higher energy level, it is its excited state.

the excited state is not permanent, it will lose the energy and return to ground state

Question 10

state the expression for the maximum number of electrons able to be housed in each shell

Answer 10

2n^2 where n is the physical quantum number/distance of electron from nucleus/shell number

ie. 2(1)^2 =2, 2(2)^2 =8, 2(3)^3 =18

Question 11

explain subshells/ sub levels

Answer 11

each shell comprised of one or more subshells specified by a secondary quantum number, l

s,p,d,f specify the value of l

each subshells energy differs: f > d > p > s

it tells us the shape of the region of space where electrons might be found

Question 12

define & explain atomic orbitals

Answer 12

(1) region of space within an electron subshell where electrons are most likely to be found
(2) the probability (heavily linked to quantum theory) of finding an electron within this space as electrons are constantly, rapidly moving

each subshell comprises of one or more atomic orbitals

Question 13

explain how to describe atomic orbitals / distribution of electron density in an atom

Answer 13

3 quantum numbers:

n, principal quantum number (or shell number/energy level), describes its size

l, angular moment quantum, denoting the sub levels (s,p,d,f), describes its shape

m, magnetic quantum number, is associated with the orientation of the orbital angular momentum

Question 14

explain shape of subshells

Answer 14

s subshells (sharp), are spherical in shape but differ in size (n) ie. 1s < 2s < 3s

p subshells (principal), consists of 3 orbitals whose directions lie at 90 degrees to one another. ie. px, py, pz

Question 15

explain pauli exclusion principle

Answer 15

states that each orbital can hold up to a maximum of 2 electrons, which must be of opposite spin.

this is how an atomic orbital is represented, by a box with either unpaired or paired electrons

Question 16

state the maximum number of electrons in a subshell

Answer 16

s - 1 orbital - max 2 electrons

p - 3 orbitals - max 6 electrons

d - 5 orbitals - max 10 electrons

f - 7 orbitals - max 14 orbitals

Question 17

state aufbau’s principle

Answer 17

electrons fill orbitals of lowest sub energy level to capacity, before filling orbitals of a higher sub energy level

aka subshell filling sequence

Question 18

explain hund’s rule

Answer 18

for orbitals of equal energy, electrons will each occupy an orbital before pairing up

ie. ‘p’ subshell has 3 orbitals of equal energy, ‘d’ subshell has 5 orbitals of equal energy, ‘f’ subshell has 7 orbitals of equal energy

Question 19

explain deviations to aufbau’s principle

Answer 19

Cr (24) = [Ar]4s13d5 NOT [Ar]4s23d4

this is because 4s and 3d subshells have almost similar energy level hence there is slightly greater stability of d subshells that are half-filled. the electrons occupy the orbitals singly to ease electron-electron repulsion.

Cu (27) = [Ar]4s13d10 NOT [Ar]4s23d9

there is slightly greater stability of d subshells that are completely filled

Question 20

explain isoelectronic species

Answer 20

any group of atoms or ions with the same number of electrons are isoelectronic

Question 21

explain how an elements position in the PT can predict its properties

Answer 21

elements of the same group tend to share similar properties