Chemistry

Matter consists of parts

that cannot be cut down

anymore, so called

„atomos“ (which means

uncuttable).

1. Matter consists of tiny particles called atoms

2. Atoms are indistructable. During reactions they

rearrange but they do not themselves break

apart.

3. In a sample of a pure element all atoms are

identical in mass and other properties.

4. Atoms of different elements differ in mass and

other properties

5. When different elements combine to compounds,

new and more complex particles are built–

however, in a given compound they are always

present in the same fixed numercial ratio

in a chemical reaction mass is neither created nor destroyed.

the individual elements that constitute a chemical compound are always present in a fixed ratio'

The „plum pudding“ model

Discovery of the electron in

1897

Positive “cloud” (the

pudding) surrounds the

negative charges (the

plums) then still called

“corpuscles“

The Gold-foil experiment

try to shoot alpha-rays

at a gold foil

There are two parts of the atom:

The core (nucleus) and the shell

core:

positive charged

massive particles

1/10.000 th of the size

protons & neutrons

shell:

negative charged

nearly mass-less particles

almost the whole atom size

electrons

His set of equations —

Maxwell's equations

— demonstrated that

electricity, magnetism

and even light are all

manifestations of the

same phenomenon:

the electromagnetic

field.

Energy of a photon of

electromagnetic radiation is

proportional to the

radiation´s frequency

E = h . ν

with h being Plank´s constant

Niels Bohr proposed a model for the hydrogen atom that incorporated quantized electron orbits. Electrons were allowed to exist only in certain orbits without emitting radiation.

Bohr introduced the idea that electrons orbit the nucleus in fixed energy levels or shells. Electrons could jump between these levels by absorbing or emitting energy in discrete quanta.

Bohr's model successfully explained the spectral lines of hydrogen but had limitations and couldn't fully account for the behavior of atoms with more than one electron.

starts to fail mathematically with

more than 1 electron

➔ Only hydrogen can be treated

with this model

So the model violates the uncertainty

principle in that it considers electrons

to have known orbits and a definite

radius, two things which can not be

directly known at once.

certain pairs of physical

properties, such as position and

momentum, cannot be

simultaneously known to

arbitrarily high precision.

Elecrtons may not be seen as a just

a particle

but also as a wave

proposed that electrons exhibit both particle and wave-like properties. This wave-particle duality was supported by experimental evidence.

The Wave Function

3 dimensional waves

Starting from the hydrogen atom the

wave function Ψ

can be used to describe the electrons

Erwin Schrödinger developed a mathematical equation (Schrödinger equation) describing the behavior of electrons as waves.

The wave function, denoted by ψ, represents the probability amplitude of finding an electron at a particular location in space and time.

The square of the wave function (|ψ|^2) gives the probability density of finding the electron.

Unlike Bohr's orbits, the wave function provides a more continuous and probabilistic description of the electron's location.

Since the location can not

be given precise

Ψ²

Probability of the

position of the electrons

these spaces are called

Orbitals

= number of the „shell“ from Bohr´s model

n = 1 to ∞

sometimes also refered to with capital

letters

n = 1 = K; n = 2 = L; n = 3 = M; ….

gives the size of the wave

= subshells

value of l depends on value of n

there are subshells from 0 to n-1

➔bigger shell contains more subshells

Ex. n=1 ➔ l=0

n=2 ➔ l=0 or 1

n=3 ➔ l= 0 or 1 or 2

= orientation of the orbitals

To each value of l there are orientations

from –l to +l

Ex.

l=1 ➔ m= -1,0,+1

l=2 ➔ m= -2,-1,0,+1,+2

That means that every orbital holds

2 electrons !

Those 2 electrons differ only in their

so called „spin“

s = - ½ or + ½

states that electrons fill lower-energy atomic orbitals before filling higher-energy ones

1) E-minimum

Orbitals of lower energy are occupied first

When there are different orbitals of the same energy they

first get occupied with a single electron

Electrons do at least differ in 1 quantum number →

In one orbital the two electrons have to differ in their spin

Size increases

from the top to

the bottom and

from the right

to the left

energy required

to remove an

electron

energy change

associated with

the adding of an

electron

(Noble gases don´t have

a positve electron

affinity!)

The ability to attract an electron

Overall:

Inreases towards fluorine

(from bottom to top and from left to

right)

the element with the highest

electronegativity

The value of ΔH° for any reaction that

can be written in steps equals the

sum of the values of ΔH° of each of

these individual steps

2 or more atoms react when during

this reaction energy is released

→they want to keep their energy

level as low as possible

→Achieved through reacting with

each other!

The s- and p-electrons are called

valence electrons or outer electrons

since they build the beginning (s) and

the end (p) of each period

So having 8 valence electrons (an octet

of electrons) means having the state of

a noble gas (configuration s2p6)

Sometimes a bond can be established

by both bond electron comming from

one partner ex.: NH4+

Equilibrium systems always

conuteracts a disturbing

influence to return to the

equilibrium state.

Any change in status quo

prompts an opposing

reaction in the responding

system.

reverse reaction → inverse K

•reaction multiplied by x → K raised to the power of x

•K for a net reaction → product of the individual steps

To predict the direction of equilibrium of a reaction, the

reaction quotient Q is used.

Q, is calculated in the same way as the equilibrium

constant, only that Q uses the current or initial

concentrations instead of the equilibrium

concentrations used to calculate K

If Q < K, then there are more reactants present than

equilibrium and the reaction will need to produce more

products to reach equilibrium and will shift to the right.

If Q > K, then there are more products present than

equilibrium and the reaction will need to produce more

reactants shifting the reaction to the left.

If Q = K, then the reaction is already at equilibrium and

there will be no shift.

Intermolecular forces between the

molecules in a phase (e.g. liquid), which

hold the phase together

Intermolecular forces between molecules of

different phases and composition (e.g. liquid

water /solid metal; liquid water / liquid

ether)

Surface tension describes the force, which holds

the liquids together. It can be quantitatively

described by the energy required to increase

the surface area of a liquid.

Viscosity describes the

gooey appearance of

liquids caused by the

inner friction between

the molecules

is the ability of two

substances to form a homogeneous

solution.

The solute is the substance that is being dissolved, while the solvent is the dissolving medium.

telített oldat

is a solution that holds as

much solute as possible.

At this state an equilibrium between the solute

in the solid, liquid or gaseous state and the

dissolved state is reached.

Saturation is temperature dependent.

Increasing or decreasing the temperature will

alter the amount of solvent needed to reach

saturation.

The solubility product constant is a measure for

the maximum solubility of a salt in water.

The smaller pKa, the stronger the acid. The smaller pKb, the stronger the base.

The higher Ka and Kb values, the stronger the acid or base.

There are three general ways charge

separation can occur in a molecule.

Formation of ions

•Permanent dipoles caused by a polar bond

•Induced dipoles in polar and non polar

molecules.

Charges at opposite end of a dipole,

that are fractions of full 1+ or 1-

charges. They are caused through an

uneven distribution of electron density

in a molecule with respect to its nuclei.

Partial charges are normally indicated

by δ+ and δ-

.

formal charge bc of coordinate bond

partial charge bc of polar bond

Valence shell electron pair repulsion

Every electron pair wants as much space

as possible

e.g.: CH4

Van der Waals Interactions

describe the electrostatic interactions

between non charged molecules,

which are caused either the formation

of a polar bond (permanent dipoles) or

by a temporary disequilibrium in

electron density across the molecule

(induced dipole).

-dipole dipole interaction

-hydrogen bond

-London Dispersion Interactions

The process of generating a cage-like network of a

solution‘s solvent molecules around a molecule or ion

of the solvent

is the special case of water being the solvent

An acid is a substance that reacts

with water to produces a

hydronium ion H3O+

A base is a substance that produces

hydroxide ions OH- in water

an acid-base reaction is a proton

transfer reaction

(the proton is the H+)

An acid needs a base to react

with

This theory doesn´t need water

to react with the acids and bases!

Oxoacids are formed through the reaction of

nonmetal oxides with water

ex:

SO3 + H2O ⇄ H2SO4

CO2 + H2O ⇄ H2CO3

the oxides are therefore also called acetic

anhydrides („without water“)

compounds that can donate or accept

a proton

acids: sour, corrosive, change litmus to red

base: slippery, change litmus to blue

acid: any ionic or molecular species that can

accept a pair of electrons to form a coordinate

bond

base: any ionic or molecular species that can

donate a pair of electrons to form a coordinate

bond

neutralisation: formation of the coordinate

bond between donator and acceptor

are weak organic acids

if deprotonate change color

color 1 as acid, color 2 as conjugated base

universal indicator: mixture of several

indicators (red-green-blue) pH: 3-7-11

phenolphtalein: colorless-purple 8.2-10

bromthymolblue: yellow-blue 6.0-7.6

methyl-orange: red-yellow 3.2-4.4

red cabbage: natural universal

indicator

electrode system

no more than a voltmeter that displays measurements

in pH units instead of volts

Buffers are watery solutions, where the

pH value stays constant upon the

addition of small amounts of

concentrated acids or bases. They are

usually composed out of an acid with

intermediate or weak strength and the

corresponding base pair.

Oxidation describes the process of taking up

oxygen or releasing hydrogen.

Reduction describes the process of releasing

oxygen or taking up hydrogen.

Oxidation takes place, whenever a compound

releases electrons. It is indicated by an increase

in the oxidation number

Reduction takes place, whenever a compound

takes up electrons. It is indicated by a decrease

in the oxidation number

big EN sum and low EN difference - covalent

small EN sum and small EN difference metallic bond

big EN difference ionic bond

A hypothetical reaction that constitutes

exclusively either the reduction or the

oxidation half of a redox reaction and in

whose equation the correct formulas for all

species taken part in the reaction are given

together with enough electrons to give the

correct electrical balance

Potential

The potential describes the ability of a compound to

perform work.

Chemical potential (μ)

The chemical potential describes the increase in the ability

of a compound to perform usable work during transport of

a certain amount of the compound.

(e.g. a concentration gradient)

Electrical potential (ϕ)

The electrical potential refers to the work, which can be

performed through charges moving in an electrical field.

(e.g. a positive charge moving to the negative pole)

To quantify differences in reactivity of redox

reactions standard reduction potentials were

established.

The standard reaction potential describes the

tendency of a reduction reaction to take up

electrons at standard conditions.

The higher the tendency to take up electrons,

the higher the potential.

The reduction potential of a half

reaction at 25°C when all ion

concentrations are 1M and the partial

pressure of all gases are 1 atm. It is

measured as a voltage difference

between the half reaction of interest

and the standard hydrogen electrode

(SHE) as a standard.