atomic structure and bonding Flashcards
atoms
atoms are basic and smallest particles that matter is made out of
compound
composed of two or more separate elements
molecule
consists of two or more atoms, can consists of any number of elements (1 or 2 or more), a group of atoms bonded together
isotopes
atoms with the same number of protons but a different number of neutrons
same electron configuration and possess similar chemical properties but have different physical properties
cations
positive ions
anions
negative ions
atomic radius: definition
size of the atom. (the distance from the centre of the nucleus to the boundary of the surrounding cloud of electrons)
atomic radius increases…
⬅️⬇️
atomic radius: ACROSS
decreases along a period as the nuclear charge increases, which lead to the increasing attraction of the electrons bringing them closer to the nucleus as the number of protons increase, hence a smaller atomic radius
atomic radius: DOWN
increases as you go down a group as the number of shells increase, hence having a greater distance between the nucleus and the outer electron shell (greater radius)
1st ionisation energy: definition
minimum amount of energy required to remove the single loosely bound outermost electron from a “gaseous” atom (in gaseous phase)
1st IE: formula
Na(g) + E1 → Na+(g) + e- (electron)
1st IE increases…
➡️⬆️
1st IE: across
ionisation energy increases as you go along a period due to the atom’s increasing nuclear charge and decreasing atomic radius, indicating that there’s a stronger attraction between the electrons and the nucleus, hence needing a higher amount of energy to remove the single loosely bound outermost electron from a “gaseous” atom
1st IE: down
ionisation energy decreases as you go down a group as the number of electron shells increases hence electrons are further from the protons’ force of attraction and less energy is required to remove the loosely bound electron as it is easier to remove.
electronegativity: definition
the atoms attraction for electrons in a bonding situation
electronegativity increases…
➡️⬆️
what elements have the highest values of electronegativity?
general non-metals have the highest values of electronegativity as they gain electrons to form ions, conversely metals have low electronegativity
Fluorine has highest electronegativity followed by O, Cl/N, C/S
electronegativity: across
across a period, values of electronegativity increase. down a group, values of electronegativity gradually decrease
smaller AR, greater attraction for electrons as the valence electrons are closer to the attractive force of protons
electronegativity: down
larger AR, outer shell is further ways from nucleus, hence lower force of attraction of protons, hence lesser attraction for electrons
increase in nuclear charge leads to increase of the shielding effect but it has minimal effect on radius going down a group
shielding
the shielding effect described the decrease in attraction between an electron and the nucleus in any atom due to an increase in number of shells
the inner shells shield the outer electrons from the attractive force of protons, ⬆️shells ⬆️shielding
this ⬇️attraction to protons and ⬇️ ionisation energy
core electrons shield the valence electrons from the full attractive forces of the protons in the nucleus
nuclear charge effect on shielding
increase in nuclear charge leads to increase of the shielding effect but it has minimal effect on radius going down a group
periodic trends
- atomic radius
- 1st ionisation energy
- electronegativity
- shielding
what are chemical bonds caused by?
chemical bonds are caused by electrostatic attractions that arise because of the sharing or transfer of electrons between participating atoms
valency
a measure of the bonding capacity of an atom
the ability of atoms to form chemical bonds can be explained by…
the arrangement of electrons in the atom and in particular by the stability of the valence electron shell
covalent molecular substances
small groups of atoms become covalently bonded to one another forming many small clusters of atoms
covalent molecular substances: examples
Carbon Dioxide CO2, Water H2O, Oxygen O2
covalent molecular substances: properties
Non-conductors of electricity in either solid, liquid or aqueous phase*
- localised electrons in a covalent molecular substances cannot freely move independently and do not contain ions, absence of freely mobile charged particles explains why covalent molecular substances are non-conductors of electricity
- * some exceptions ⬇️⬇️
Some are good conductors of electricity when in aqueous phase
- acidic/basic substances react with water and ionises producing free mobile ions which can move freely throughout the solution carrying a charge hence conducting a current
Soft and weak
- intramolecular forces (strong) forces (between the atoms within the molecules), weak intermolecular forces between neighbouring molecules, hence molecules can easily be separated
Low to moderate melting and boiling points
- weak intermolecular forces easily broken, weakly bonded lattice which can be easily disrupted by heat energy, intramolecular forces are unaffected
covalent network: structure
Within a covalent network material atoms are covalently bonded to one another forming a single vast array that involves every atom within the sample
difference between covalent molecular and covalent network
The vast difference in properties between these two is caused by the difference in structure.
covalent network: examples
Silicon (Si), Silicon Dioxide (e.g. quartz, SiO2), Diamond and Graphite (C)*
*elemental carbon exists as a range of allotropes, including graphite, diamond and fullerenes, with significantly different structures and physical properties
covalent network: properties
Non-conductors of heat or electricity *
- electrons are in fixed positions within the atom’s shell, cannot freely move hence unable to conduct electricity and carry heat through substance
- GRAPHITE: some electrons are delocalised hence can move freely through the substance
Very hard and brittle *
- difficult to disrupt the array of strongly covalently bonded atoms
- GRAPHITE: has weak covalent bonds allowing layers to slip over one another
Very high melting and boiling points
- strong covalent bonds occur between all the atoms within the structure, hard to disrupt this continuous array of strongly bonded atoms
*GRAPHITE IS AN EXCEPTION
metallic substances: structure
metal atoms lose their valence electrons and form positive metal ions and free electrons. the resulting metal ions occupy fixed positions forming a regular three dimensional lattice. valence electrons released by the metal atoms move freely amongst the positive metal ions. these free moving electrons are sometimes are sometimes referred to as a sea of electrons.
stability of the metallic structure
the stability of the metallic structure comes from the strong electrostatic attraction between the positive metal ions and the sea of electrons. this electrostatic bond, known as the metallic bonds acts in all directions between the positive metal ions and the sea of electrons and so is referred to as a non-directional bond.
metallic substances: examples
Copper (Cu), Iron (Fe), Aluminium (Al)
metallic substances: properties
high melting and boiling points
- Strong forces of attraction between positive metal cations, fixed positions in metallic lattice, delocalised electrons.
good conductors of electricity
- valence electrons are mobile, freely moving, carries a charge, conducts electricity or a current.
good conductors of heat
- electrons freely moving and can carry heat energy as kinetic energy, can carry a flow of heat through metallic lattice
malleable and ductile
- non-directional bonding between metal ions and the Sea of delocalised electrons and can move allowing metal to change shape without disrupting the bonding.
lustrous (shiny)
- light rays reflect off delocalised electrons demonstrating that metal shine, preventing light to pass through, metal is opaque.
ionic bonding
ionic bonding can be modelled as a regular arrangement of positively and negatively charged ions in a crystalline lattice with electrostatic forces of attraction between oppositely charged ions
ionic substances: examples
Magnesium Iodide (MgI2), Sodium Chloride (NaCl)
ionic substances: properties
poor conductors of electricity in the solid phase
- charged particles (ions) are in fixed positions in lattice hence being unable to move and carry charge. Electrons within the ionic lattice are also tightly held by individual ions and hence also unable to move and carriage charge through the ionic solid
good conductors of electricity in the molten phase and when dissolved in water (aqueous phase)
- In the molten and aqueous phase both positive and negative ions are mobile and conduct an electric current as they can move independently of one another. the positive ions move to the negative electrode and the negative ions move towards the positive electrode.
hard and brittle
- if a large force is applied to an ionic lattice it will cause layers of ions to move, like charges will be closer to each other due to this, the repulsive forces between the like charged particles will cause lattice to break
high melting and boiling points
- Ionic Bonds which are strong electrostatic attractive forces between ions extend throughout the ionic lattice keeping individual ions in fixed positions hence a high temperature is needed to disrupt the ionic lattice
theories of atom development: people
Due To Red Buses Crashing
D - John Dalton
T - J.J. Thomson
R - Ernest Rutherford
B - Neils Bohr
C - James Chadwick
john dalton
precise definition of the indivisible building blocks of matter that we now know as atoms. His atomic theory can be summarised as follows:
- All matter, whether an element, a compound or a mixture is made out of very tiny particles called atoms which are indivisible and indestructible, elements are composed of extremely small particles called atoms
- All atoms of a given element are identical having the same size, mass and chemical properties. Atoms of different elements have a different size, mass and chemical properties. size, and chemical properties
- Atoms are not created nor destroyed or changed into different types during a chemical reaction
- a chemical reaction involves only separation, combination or rearrangement of atoms
- Compounds are formed when atoms of more than one element combine in a specific ratio, The relative number and kinds of atoms are constant in a given compound.
j.j. thomson
J.J. Thomson discovered the electron by experimenting with a Crookes or a Cathode ray tube. He proposed a model of the atom which he likened to a plum pudding. The negative electrons represented the raisins in the pudding and the dough contained the positive charge. Atoms are matter with negatively charged electrons and better in the positive substance.
Ernest Rutherford
beam of alpha particles were targeted at a very thin sheet of gold foil only a few atoms thick. Rutherford discovered the nucleus using the gold foil experiment. He demonstrated that the atom has a tiny and heavy nucleus and that most of the atom is empty space.
Neils Bohr
Bohr proposed his quantised shell model of the atom to explain how electrons can have stable orbits around the nucleus. Electrons should move around the nucleus but only in prescribed orbits. When jumping from one orbit to another with lower energy, a light quantum is emitted.
james chadwick
Chadwick is best known for his discovery of the Neutron. Chadwick bombarded Beryllium with Alpha particles. An unknown radiation was produced and this particle became known as the Neutron.
subshells
s - subshell: 1 orbital
- can hold 2 electrons
p - subshell: 3 orbitals
- can hold 6 electrons
d - subshell: 5 orbitals
- can hold 10 electrons
f- subshell: 7 orbitals
- can hold 14 electrons
pure substances
a pure substance consists only of one element or one compound
mixtures
a mixture consists of two or more different substances, not chemically joined together
Homogenous mixture
no compounds, mixture of elements, usually solutions
Heterogenous mixture
mixture of compounds
flame tests
- When atoms are heated their electron(s) jumps from its ground state to the excited state.
- When the atom cools the excited electrons(s) go back to their ground state and emit the absorbed energy as a specific wavelength that we see as a certain colour.
- Different elements produce different colours as they have different electron configurations and their electrons jump to different shells.
- The distance(s)/differences in energy between ground state and excited state is different for each atom and they therefore absorb different amounts of energy.
- They therefore emit different amounts of energy when they go back to the ground state and as a result they produce different colours.
separation techniques
filtration and evaporation are techniques that can be used to separate the substances in a mixture based on the differing solubility and volatility
distillation can be used to separate a mixture based on boiling point differences in the individual components of the mixture
diamond
in diamond, each Carbon atom is bonded tetrahedrally to four other carbon atoms, the carbon atoms are held together by strong carbon-carbon single covalent bonds
graphite
in graphite, the carbon atoms form covalently bonded layers that are held together by intermolecular forces, each carbon is covalently bonded to three other carbons in the same layer to form interconnected hexagonal rings, there are delocalised electrons that bounce between the hexagonal layers
layers can separate the structure easily
