Atomic StructureIB Chemistry Topic 2

The Atom

Originally thought to be the smallest form of matter due to the fact that the atom can’t be broken down into simpler components by chemical reaction.

Dalton’s Atom• All matter is composed of tiny indestructible particles called atoms

• Atoms cannot be created or destroyed

• Atoms of the same element are alike in every way

• Atoms of different elements are different

• Atoms can combine together in small numbers to form molecules (compounds)

Dalton’s Atom• Dalton’s Law of Constant Composition is seen in the chemical formulas of compounds

• Each atom is represented by its element symbol

• The number of each type of atom is indicated by a subscript written to the right of the symbol

• H2O N2O5 C6H12O6 Fe2(CO3)3

J.J. Thompson• J.J. Thomson (1856-1940) –atoms contain negative particles (electrons)

• Proposed the “plum pudding” model of the atom - if a negative charge was present in atoms, there must also be a positive to balance it; negative charges were distributed evenly on a positive sphere (raisins suspended on pudding)

Gold Foil Experiment• Ernest Rutherford (1871-1937) aimed positively charged particles at a thin metal foil; most particles went through but some were deflected in different directions

• (a)If the plum pudding model was correct.

• (b)Actual results

Structure of the Atom• Protons + neutrons = nucleons located in the

nucleus• Electrons are found in energy levels or shells

surrounding the nucleus

• Most of the atom is empty space

Structure of the Atom

Describing the atom• Mass Number (A) = number of protons + number of neutrons in an element

XA

Z

n+ / n-Charge if it is an ion, zero if not

• Atomic Number (Z) = the number of protons (also electrons in a neutral atom). The atomic number is unchanging, it is the unique identifier of an atom.

Label A and Z for the following:

Be Ge

O Te

Mo Rb

Isotopes• Atoms of the same element (same atomic #) with different masses due to a different number of neutrons .

• Examples: 1H, 2H, 3H C-12, C-14 35Cl, 37Cl

Deduce the symbol for an isotope given its mass number and atomic number.

1. p = 6

n = 7

2. p = 3

n = 4

3. p = 16

n = 17

e = 18

Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge.

1. Cle=?

p=?

n=?

2. Which subatomic particle occurs in the same amount in both of the following species?

P S

35

17

31

153216

2-

Properties of Isotopes• All isotopes of an element have the same chemical properties because chemical properties are determined by the protons and electrons, not the mass.

• Isotopes have different physical properties.

Ex: rate of diffusion, mass, density, melting point, boiling point.

Radioisotopes

•Some isotopes are radioactive (the nuclei of these atoms break down spontaneously and emit radiation).

Radioisotopes•Uses:•Nuclear power generation, sterilization of surgical instruments in hospitals, crime detection, finding cracks and stresses in structural materials, food preservation…

•Add to your notes: Iodine – 131, Cobalt – 60, Carbon – 14, (brief description and uses) •Shroud of Turin•Pgs. 44-46

The electromagnetic spectrum

• Visible light only makes up a small part of the electromagnetic spectrum.

𝐸=h𝜐=h𝑐𝜆

Different types of spectraA continuous spectrum is produced when white light is passed through a prism and shows all the frequencies.

Line spectra • When white light passes through

gases certain absorption will occur • Results in a line spectrum produce

with some colors of continuous spectrum missing

• Different elements have different line spectra

• Therefore elements can be identified by their line spectra similar to products with their barcodes

Transitions and EMS

Series nf ni Region of EMS

Lyman 1 2,3,4,5… UV

Balmer 2 3,4,5,6… Visible and UV

Paschen 3 4,5,6,7… IR

You don’t have to knowthe series name, but youdo need to know whichregion of the EMS goes with which transitions.

Hydrogen spectrum:The type of radiation given out by an atom is dependent on where an electron falls to from its excited state

Balmer:

Heisenberg’s uncertainty principle: It is impossible to accurately know both the position and the momentum of an electron. The more we know about the position of an electron, the less we know about the momentum, and vice versa.

Schrödinger’s equation:A very complex math equation whose results are describes by atomic orbitals. The results describe the probability density of the space an electron can occupy.

An atomic orbital is a region in space where there is a high probability of finding an electron.

Orbitals• Energy shells are divided into 4 levels:

Principal quantum number, n, has integer values (1, 2, 3…) and can hold 2n2 electrons. This is the main energy level.

Sublevels, l: s, p, d, and f. Each of these levels can hold 2 electrons, but the number of orbitals, ml, in each sublevel is different:

Sublevel # of orbitals Max e- number

s 1 2

p 3 6

d 5 10

f 7 14

Orbital Shapes

Orbitals• Orbitals in sublevel p are labeled px, py, and pz.

**Labels for d and f do not need to be known.

• Within each orbital, are two electrons that have spins, ms, with values of +1/2 or -1/2.

• Pauli exclusion principle: an orbital can hold two electron and they must have opposite spins.

• Hund’s rule: electrons will fill each degenerate orbital (orbitals of equal energy) singly, before occupying them pairs.

Orbital DiagramsWhen drawing orbital diagrams, we use the arrows-and-boxes method.

For each sublevel, there is one box per orbital.

For s: 1 boxFor p: 3 boxesFor d: 5 boxesFor f: 7 boxes

Orbital diagramsAufbau principle: electrons fill the lowest-energy orbital that is available first.

Ex: Sulfur

16 electrons

Electron Configuration• Full electron configuration:

Sulfur: 1s12s22p63s23p4

• Electron configuration shorthand: Sulfur: [Ne]3s23p4

Remember this from CP?