Please enable JavaScript.
Coggle requires JavaScript to display documents.
Chem Chapter 1: Atomic Structure and physical periodicity - Coggle Diagram
Chem Chapter 1: Atomic Structure and physical periodicity
subatomic particles of atom
proton
relative charge: +1 relative mass: 1
electron
relative charge: -1 relative mass: 1/1836
neutron
relative charge: 0 relative mass: 1
ions and deflection of charged particles
a negative ion has more elections than protons; known as anion
positive on has fewer negative electrons than protons; known as cation
protons are deflected to the negative plate, whereas electrons are deflected to the positive plate. electrons are deflected to a much greater extent as it has a lower mass than proton
ANGLE OF DEFLECTION IS DIRECTLY PROPORTIONAL TO CHARGE, AND INVERSELY PROPORTIONAL TO MASS
RELATIVE ATOMIC MASS: average mass of one atom of a element/ 1/12 mass of carbon-12 atom
RELATIVE ISOTOPIC MASS: mass of one atom of an isotope of an element/ 1/12 mass of carbon-12 atom
arrangement of electrons
principal quantum shell: describes main energy level of electron
also indicates relative size of orbital and relative distance of an electron from the nucleus
larger value of quantum number, the higher the energy level and further the electron is from the nucleus
maximum number of electrons that can occupy one quantum shell is given by formula 2n^2
sub shell: each principal quantum shell can be divided into one or more subshell
no. of subshell is the same as principal quantum number
MAX no. of electrons in a orbital is : TWO
each subshell has one or more orbitals with the same energy
level but different orientations in spaces
S ORBITAL: all S orbitals are spherical in shape,
size increases with the principal quantum number
P ORBITAL: p sub shell has 3 p orbitals
each orbital can hold up to 2 electrons, p sub shell can hold up to 6 electrons
dumbell in shape
3 p orbitals are identical in size but differ in their spatial oritntaiton
represented as px, py and pz
D ORBITAL: d subshell has 5 orbitals
each d subshell can hold up to 10 electrons
the 5 d orbitals are also degenerate (have the same energy level)
electronic configuration
guidelines: fill in all the individual boxes before pairing
exceptions are Cr and Cu (fill in all the 3d electrons before the 4s)
when writing for cation, write the full electronic configuration of the atom first before subtracting the electrons
same thing for anion, but add the electrons
electronic configuration in periodic table:
GROUP
eg valence shell electronic configuration in group 1: ns1
eg " " " " " " 2: ns2
group 3: ns2 p1
PERIOD
eg same shells
periodic trends in atomic and ionic radii
IMPT: HOW TO ANSWER:
number of quantum shells; determine how far the outermost electron is from the nucleus, further the distance, weaker the attraction between the outermost electron and nucleus
effective nuclear charge: Zeff, met clear charge experienced by an outer electron, must do so by comparing nuclear charge and shielding or screening effect
nucelar charge
dependent on the number of protons
protons, > nuclear charge
shielding effect
presence of inner shell electrons reduce the electrostatic attraction between the outermost electrons and the nucleus
HOWEVER shielding effects small compared to that by the inner core electrons
ATOMIC RADIUS
across period 2 and 3: atomic radius DECREASES down the group
nuclear charge increases due to increasing number of protons
shielding effect remains relatively constant
effective nuclear charge increases, reusltng in stronger electrostatic forces of attraction between nucleus and outermost electron
outermost electrons are oiled closer to the nucleus and hence a decrease in atomic radius
across first row transition elements: atomic radius is relatively invariant
increasing proton charge is nullified by increasing shielding effect as electrons start to fill up the 3d subshell
effective nuclear charge remains almost constant
Down the group: atomic radius INCREASES down the group
quantum shells increases, outermost electrons are further away from the nucleus, hence atomic radius increase.
both nuclear charge and shielding effect increase down the group, hence effective nuclear charge differs little down the group
IONIC RADIUS
ionic radius of cations and anions decrease across period 3 - however, there is sharp increase in ionic radius from cations to anions, as anions have one more quantum shell of electron than cations
ionisation energy
1st ionisation energy
: energy required to remove one mole of electron from one mole of gaseous atoms to form 1 mole of unpositively charged gaseous ions
TREND IN FIRST IONISATION ENERGY
across periods 2 and 3
increases across a period, nuclear charge increase, while shielding effect remains the same, Zeff increases and the electrostatic FOA between outermost electrons and nucleus become stronger, more energy is required to remove the outermost electron
exception
between group 2 and 13,
Al has lower 1st ionisation energy than Mg
the 3p sub shell of al is further away from the nucleus than 3s subshell.
weaker attraction between nucleus and outermost electron. hence less energy is required to remove 3p electron from al, resulting in a lower ionisation energy for al
exception between group 15 and group 16 elements between P and S
3p electrons In p are UNPAIRED whereas 3p electrons in S are PAIRED.
some inter electronic repulsion between paired electrons in the 3p subshell in s
less energy required to remove one of the paired electrons in S
across transition elements
relatively invariant
down the group
decrease as number of quantum shells increases... etc - both nuclear charge and shielding effect increases, hence Zeff does not differ emuch
2nd ionisation energy
: energy required to remove 1 mole of electrons form 1 mole of unpositivley charged gaseous ions to form 1 mole of gaseous ions with double positive charge
factors affecting ionisation energy of an atom:
number of quantum shells: larger the number, the further the electron is away form the nucleus, easier to remove, lesser ionisation energy
effective nucelar charge Zeff: higher Zeff, stronger attractive forces between nucleus and electrons to be removed, hence greater ionisation energy
ionisation energy always increases from first to second to third etc. this is because when an electron is removed from a neutral atom, the number of protons exerting attraction for the remaining electrons remains the same.
shielding effect is now reduced as their is now one less electron, hence effective nuclear charge increases,
more energy is needed to remove another electron from the more positively charged ion, hence ionisation energy increases