IB Chemistry: DP1 S1 - The atom, Periodicity, Bonding

Who first proposed the periodic table

Mendeleyev in 1869 by grouping elements into families, leaving gaps for elements that exsist but have not yet been discovered

Metals

●Good conductors of heat and electricity

●Malleable - can be beaten into thin sheets

●Ductile - can be drawn into wires

●Shiny or have luster

●Tend to lose electrons in chemical reactions

Non-metals

●Poor conductors of heat and electricity

●Tend to gain electrons in chemical reactions

●Can be solid, liquid or gas

Metalloids

have both metallic and non-metallic properties

Group 1 - Alkali metals

silvery metals that are too reactive to be found in nature - they are stored in oil to prevent contact with air and water.

The first 3 elements have the following properties:

●Physical

○Good conductors of electricity

○Low densities

○Have grey shiny surfaces when cut with a knife

●Chemical

○They are very reactive metals.

○They form ionic compounds with non-metals.

They form single charged ions with a 1 plus charge. (M+)

-Reactivity increases down a group because ionization energy decreases and the outer shell electron is more easily removed.

-Their ability to conduct electricity is due to the mobility of the outer shell electron

Group 17 - Halogens

exist as diatomic elements.

●Physical properties:

○They are colored.

○They show a gradual change from gases (F2 and Cl2) to liquid (Br2) to solids (I2 and At2).

●Chemical properties:

○They are very reactive non-metals.

○Reactivity decreases down the group.

○They form ionic compounds with metals or covalent compounds with other non metals.

-Reactivity decreases down the group as the atomic radii increases and the attraction for outer electrons decreases.

Group 18 - Noble gases

●contains the least reactive elements.

●With the exception of Helium, they have complete outer shells of 8 electrons: a stable octet.

●Their lack of reactivity is explained by their inability to lose or gain electrons.

-Noble gases - are colorless gases, exist as single atoms (monatomic), and are very unreactive

Group 1 reactivity with water

-Li reacts slowly, releases H2 but keeps its shape and floats.

-Na reacts with a vigorous release of H2 which releases enough heat to melt the remaining Na that floats on the surface.

-K reacts vigorously to produce enough heat to ignite the H2 produced and creates a lilac flame which dances on the surface

StartFragment

Li reacts slowly, releases H2 but keeps its shape and floats.

Na reacts with a vigorous release of H2 which releases enough heat to melt the remaining Na that floats on the surface.

K reacts vigorously to produce enough heat to ignite the H2 produced and creates a lilac flame which dances on the surface

Alkali metals react with water to form hydrogen gas and the metal hydroxide.

○They are called alkali metals because they create a basic or alkaline solution when reacted with water.

DOWN THE GROUP: Reaction with water becomes more vigorous as you descend a group because the ionization energy decreases and positive ions are formed more readily

Halogens react with group 1 metals

●Halogens react with Group 1 alkali metals to form ionic or metallic halides.

●The metal loses its electron to the halide to form a halide ion X- and a positive ion M+.

●After the electron is transferred, the ions are pulled together by the strong attraction between the positive and negative charges.

●The most vigorous reactions occur between elements that are farthest apart on the periodic table. (Bottom left with top right.)

Displacement reactions of halogens

●The more reactive halogen displaces the ions of the less reactive halogens. F>Cl>Br>I

●This is seen in single replacement reactions.

The smaller the halogen, the more reactive it is.

Color changes indicate displacement reactions.

When Br2 is displaced by chlorine, the solution turns from clear to orange.

When I2 is displaced by bromine, the color darkens even more. It turns purple when mixed with a hydrocarbon.

Halogens also form insoluble salts with silver which can help identify the halide.

Acid-base character of period 3 oxides

●Metallic elements which form ionic oxides are basic.

●Non-metallic elements which form covalent or non-metallic oxides are acidic.

●Na and Mg oxides are basic.

●Al and Si oxides are amphoteric - show both acidic and basic properties.

●P, S, and Cl oxides are acidic.

Metallic character

●increases down a group and decreases from left to right or across a period.

●Metals tend to have low ionization energies as they tend to lose electrons in a chemical reaction.

●Non-metals have high electron affinities as they tend to gain electrons in chemical reactions

Chemical properties of elements according to groups

-elements in Groups 1 to 3 lose electrons to adopt the arrangement of the nearest noble gas with a lower atomic number. They are usually metals.

-elements in Groups 15 to 17 gain electrons to adopt the electron arrangement of the nearest noble gas to their right. They are usually non-metals.

-the metalloids show intermediate properties

Acid rain

any precipitation (rain, fog, mist, snow) that is more acidic than normal (pH of less than 5.6. pH below 7 is acidic)

Cause of acid rain

•atmospheric pollution from acidic gases such as sulphur dioxide and oxides of nitrogen emitted from the burning of fossil fuels.

•It is also recognized that acidic smog, fog, mist, move out of the atmosphere and settle on dust particles which in turn accumulate on vegetation as acid depositions.

•When rain falls, the acid from these depositions leak and form acid dews

Periodicity

trends in properties of elements across a period and down a group

Nuclear charge

charge of all the protons in the nucleus. It has the same value as the atomic number. The nuclear charge increases you go across the periodic table

Effective nuclear charge

net positive charge experienced by valence electrons. It can be approximated by the equation: Zeff = Z - S, where Z is the atomic number and S is the number of shielding electrons. ALWAYS less than the nuclear charge

Effective nuclear charge helps explain...

trends in chemical and physical properties.

-The outer shell electrons which determine many of the chemical and physical properties do not experience the full attraction of the nucleus because they are shielded from the nucleus and repelled by the inner shell electrons

Effective nuclear charge patterns

DOWN A GROUP: the effective nuclear charge remains about the same because the inner shell electrons "shield" the outer electrons from the nucleus. Example: Sodium has 11 electrons with 1 e- in the outer shell. The outer shell electron is shielded from the 11 protons by the 10 interior electrons in the first and second energy levels.

ACROSS A PERIOD: the effective nuclear charge increases because as you are adding protons, you are still in the same energy level so the inner electrons are not increasing

Shielding

electrons in an atom can shield each other from the pull of the nucleus

Shielding effect

describes the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell

-the more electron shells there are, the greater the shielding effect experienced by the outermost electrons

Atomic radius

half the distance between neighboring nuclei

-the atom does not have sharp boundaries so it is difficult to measure its size.

Atomic radius patterns

DOWN A GROUP: increases as you go down a Group because you are adding energy levels.

ACROSS A PERIOD: decreases across a Period because the effective nuclear charge increases which increases the attraction between the nucleus and the outer shell electrons

Ionic radius patterns

ACROSS A PERIOD: The ionic radius decreases for metals forming cations, as the metals lose their outer electron orbitals.

The ionic radius increases for nonmetals as the effective nuclear charge decreases due to the number of electrons exceeding the number of protons

DOWN A GROUP: The ionic radii increases as you move down a Group due to increasing energy levels

First ionisation energy

the energy required to remove one mole of electrons from one mole of gaseous atoms.

-Measure of energy required to remove 1 electron

Ionisation energy patterns across a period

increase across a period due to increasing nuclear charge which makes it more difficult to remove the outer shell electron.

Ionisation energy trends

●Generally there is an increase in ionization energies as one moves from left to right on the periodic table.

●As you move from left to right, the nuclear charge increases therefore making it more difficult to remove an electron.

As you move down the periodic table, there is a decrease in ionization energies because the outermost electrons are further away from the nucleus

Ionisation energy patterns down a group

decrease down a group. This is due to the distance from the nucleus and the shielding by the inner shell electrons. These 2 factors make it easier to remove the outer shell electrons the farther they are away from the nucleus

Successive ionisation energy

the energy needed to remove 1 mole of each subsequent electron from each ion in 1 mole of positively charged gaseous ions.

Ionisation energy decrease group 2-3

●One would expect an increase from Be to B since they are in the same energy level, however, there is a decrease and this is due to electron configuration.

○The outer electrons being removed from Be and Mg in Group 2 are in the "s" orbitals.

○The outer electrons being removed from B and Al in Group 3 are in the "p" orbitals.

The "p" orbital electrons take less energy to remove than the "s" orbital electrons which are closer to the nucleus

Ionisation energy decrease group 5-6

○One would expect an increase from Group 5 to Group 6, but this is not so.

○Look again at the e- configurations.

■Group 5 elements have the basic configuration of ns2, npx1, npy1, npz1

■Group 6 elements have the basic configuration of ns2, npx2, npy1, npz1

An electron is easier to remove from a doubly occupied orbital because it is being repelled by its partner electron and takes less energy to pull off than electrons in singly occupied orbitals

Ionisation energy exceptions

group 16 elements have lower energies than Group 15 elements also due to electron configuration

What causes the ionisation energy exception

The Group 16 elements have a doubly occupied "p" orbital, unlike the Group 15 elements which have three singly occupied "p"orbitals.

It is easier to pull an electron from a doubly occupied orbital because of the added electron repulsion

Remember Group 3 elements have lower energies than Group 2 elements because of electron configuration.

The Group 3 elements have the general configuration of ns2, np1 and Group 2 elements are ns2. It is easier to remove the "p" electron than the "s" electrons

Electron affinity

energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions.

Also defined as the amount of energy released when an electron is added to a neutral atom or molecule in the gaseous state to form a negative ion (exothermic process hence values are negative)

Electron affinity patterns across the period

values generally get more negative as you go from left to right in a period which means that the halogens really want to add that extra electron.

Electron affinity patterns down a group

the patterns vary down a group and do not show clear trendsE

Electronegativity

measure of the tendency of an atom in a molecule to attract a shared pair of electrons towards itself (a measure of the ability of its atoms/nucleus to attract electrons in a covalent bond).

-an element with a high electronegativity has strong electron pulling power.

An element with a low electronegativity has weak electron pulling power

Electronegativity patterns across a period

increases, due to increasing effective nuclear charge

Electronegativity patterns down a group

decreases due to outer shell electrons being farther away from the nucleus which decreases the attraction

Pauling Scale

the unit used to measure electronegativity

Physical properties of transition metals

●High electrical and thermal conductivity

●High melting point (muchhigher than those of the s-block elements. The high melting points suggest that the 4s and 3d electrons are involved in metallic bonding.)

●Malleable - easily beaten into shape

●High tensile strength - can hold large loads without breaking

●Ductile - easily drawn into wires

These properties are explained by strong metallic bonding. The 3d and 4s electrons are close in energy and are all part of the delocalized sea of electrons which holds the metal lattice together

Chemical properties of transition metals

●With the exception of Zn, the 3d elements are transition metals.

○They form compounds with more than one oxidation number (Fe3+ and Fe2+)

○They form a variety of complex ions (ligands)

○They form colored compounds.

○They act as catalysts when either elements or compounds (Haber process)

They have magnetic properties (paramagnetism and diamagnetism)

Magnetic properties of transition metals

●Every spinning electron in an atom or molecule can behave as a tiny magnet.

●Electrons with opposite spins have opposing orientation so have no net magnetic effect.

●Elements and ions with paired electrons do not show magnetic properties.

●If elements or ions have unpaired electrons, they will show magnetic properties.

Why are transition element complexes coloured?

due to the absorption of light when an electron is promoted between the orbitals in the split d-sublevels. The colour absorbed is complementary to the colour observed

What is responsible for transition elements characteristic properties?

incomplete d sublevels

What explains the formation of variable oxidation states in transition elements?

the fact that their successive ionisation energies are close in value

Catalyst

a substance which alters the rate of a reaction by providing an alternative reaction pathway with a lower activation energy.

Homogenous catalysts

the catalyst is in the same phase as the reactants

Heterogenous catalysts

catalyst in different phase than reactant

What makes transition elements good homogenous catalysts?

ability of transition metals to show variable oxidation states allows them to be very effective in redox reactions

What makes transition elements good heterogenous catalysts?

ability of transition elements to use the 3d and 4s electrons to form weak bonds to small reactant molecules makes them effective heterogeneous catalysts as they provide a surface for the reactant molecules to come together with the correct orientation

What is the reason for transition metal variable oxidation numbers

because the 3d and 4s orbitals are close in energy, the electrons can be removed without a huge jump in energy as you would see from the s and p orbitals.

Transition metal variable oxidation number examples

Calcium will lose the 4s2 electrons and then it would take a huge amount of energy to pull off the electrons in the 3p orbital.

Titanium will lose the 4s2 electrons to make Ti+2, then one of the 3d electrons to make Ti+3, then the other 3d electron to make Ti+4. It does not make a +5 ion because it takes too much energy to pull off the p electrons

+2 Oxidation state

transition elements show an oxidation state of +2 when the "s" electrons are removed

-when the first row d-block elements form ions, they ALWAYS lose the 4s electrons first to make the 2+ ions.

-to make ions of higher than 2+, they start losing the 3d electrons

Why is Zinc not conisidered a transition element?

it does not form ions with incomplete d-orbitals

-Transition elements form one or more ions with a partially filled d sub-level.

-Zinc only forms one ion and it does NOT have a partially filled d sub-level.

-Zn makes the Zn2+ ion which has the electron configuration of [Ar]3d10.

-Zinc does not make colored compounds

Regular covalent bond

each atom contributes one electron

Coordinate covalent bond

a covalent bond in which one atom contributes both bonding electrons

Ligand

-molecules or ions that form coordinate bonds to the metal

-aLWAYS have at least 1 lone pair.

-can be polar molecules or anions.

Complex ion

central metal ion bonded to a group of ligands

Coordination number

the number of coordinate bonds from the ligands to the central ion

Formulas of complex ions

-metal center written first.

-parentheses used for molecule ligands (H2O, NH3, CO, CN).

-no parenthesis for single-atom ligands (F-, Cl-, Br-, etc.)

-entire complex ion is enclosed in BRACKETS, with charge outside.

-charge of complex ion = metal's oxidation state + charge on all ligands

Counterions

ions outside the complex ion, which keep the compound neutral

Polydentate ligands

ligands with two or more donor atoms

Types of ligands

monodentate, bidentate, polydentate

Monodentate

ligand binding to the centre through only one atom

Bidendate

have two donor atoms which allow them to bind to a central metal atom or ion at two points

Polydentate

range in the number of atoms used to bond to a central metal atom or ion. Ethylenediaminetetraacetate (EDTA) has six donor atoms!

Oxidation number

Positive or negative number that indicates how many electrons an atom has gained, lost, or shared to become stable

Why is anything coloured?

it absorbs visible light

-a substance has a particular color if it reflects light of that color and absorbs all others (opaque)

-it absorbs light of its complementary color (translucent) and transmits all others

Why do substances absorb visible light?

when that radiation possess the energy needed to move an electron from its lowest energy state to some excited state

Color of the complex depends on...

-the identity of ligands.

-the identity of the metal (number of d orbital electrons)

-the oxidation number of the central ion (ionic charge).

-the shape of the complex ion.

Crystal field theory

model of bonding in transition-metal complexes that accounts for many of their observed properties

•In complex ions not all 5 d-orbitals have the same energy.

•The d sublevel is split into 2 sets

-3 lower due to less repulsion

-2 higher due to more repulsion

changing the metal or the ligands causes a different energy gap between the split d orbitals.

Colorimeter

using an internal light source, a colorimeter shines light down onto the surface of the sample. As the light reflects up to the device, it passes through three filters: red, green and blue. These filters distill tri-stimulus (RGB) values that match how our eyes see color

Spectrophotometer

works almost same way, except for one main difference - the filters. Instead of using three filters to determine the RGB values of the color like a colorimeter, modern day spectrophotometers typically have 31 filters to measure the full color spectrum.

Kinetic molecular theory

model to explain physical properties of matter

Temperature

measure of average kinetic energy of particles

Kelvin conversion

T in celsius + 273.15

Element

substance that cannot be broken down into simpler substances by a chemical reaction

Atom

smallest particle (species) of an element that retains the properties of that element

Compound

the chemical combination of two or more elements

Mixtures

contain more than one element or compound in no fixed ratio, which are not chemically bonded and so can be separated by physical methods

Mass spectrometer

is used to separate difference atoms and determine their relative atomic masses and percentages

Isotope calculations problem types

-calculate relative atomic mass (give %)

-calculate % abundance

Calculate relative atomic mass (given %)

mass = (% x atomic mass) + (% x atomic mass) +....etc

Note % must be represented as decimal (70% = 0.70)EndFragment

Calculate % Abundance

mass = (atomic mass) (x)+ (atomic mass) (1-x)

Isotope

an atom with the same number of protons and a different number of neutrons from other atoms of the same element.

Chemical properties of the isotopes of an element

show the same chemical properties as their parent element since neutrons do not affect how they react

Physical properties of the isotopes of an element

Isotopes with more neutrons are heavier and move more slowly at a given temperature. This can be used to separate them.

-the difference in neutrons does affect physical properties like boiling and melting points, mass, density and rate of diffusion for gases.

-remember a physical property is something that can be measured without changing the chemical composition of the substanceEndFragment

Models of the atom

•John Dalton - Dalton's atomic theory

•JJ Thomson - plum pudding model and discovered the electron

•Ernest Rutherford - gold foil experiment and discovered the proton

•Niels Bohr - solar system model where the electrons orbit the nucleus

•Quantum Mechanical Model - modern theory where electrons exist in cloud shapes or "orbitals"EndFragment

Energy levels

the possible energies that electrons in an atom can have

Energy levels analogy

•the steps on a ladder:

-You can't stand between the steps on a ladder, and electrons cannot hang out between energy levels.

Number the energy levels: n = 1, 2, 3, 4, ...

Energy levels are different from the steps on a ladder because they are NOT evenly spaced!

Quantum mechanical model

-the modern description of the behavior of electrons in atoms

-uses a "cloud" model to describe where the electron is likely to be found

- it's based on probabilities (chances) that something will be true.

-these clouds take on shapes based on where an electron with a specific energy is most likely to be found

Emission spectra

produced by atoms emitting photons when electrons in excited states return to lower energy levels

Desctibe the electromagnetic spectrum

-electromagnetic radiation comes in different forms

-all forms travel at the same speed of light but have different wavelengths

-the higher energy forms have, shorter wavelengths and higher freuencies

Electromagnetic spectrum lowest energy to highest

-radio waves, microwaves, infrared radiation, visible light (ROYGBIV), ultraviolet radiation, x rays, gamma rays

Continuos spectrum

-a continuous spectrum is produced when white light is passed through a prism.

-it shows all colors in an unbroken sequence of frequencies, such as the spectrum of visible light.

Line spectrum

-an emission spectrum that has sharp lines produced by specific frequencies of light.

-it is produced by excited atoms and ions as they fall back to a lower energy level.

-different elements have different line spectra so they can be used to identify unknown elements

What does the line emission spectrum of hydrogen provide evidence for?

the existence of electrons in discrete energy levels which converge at higher energies

The hydrogen spectrum

-complex, comprising more than the three lines visible to the naked eye. It is possible to detect patterns of lines in both the ultraviolet and infrared regions of the spectrum as well.

-when an electron drops down from one energy level to another, it emits a very specific wavelength of light (i.e., it emits a photon with a specific energy). The farther the drop, the shorter the wavelength and the higher the energy of the photon

-when the energy levels are drawn around the nucleus, they are spaced farther apart near the nucleus and much closer together at higher energy levels