Elements & Periodic Table - A0S 1 - Chapter 2

Remembering - Definitions

Atom

Smallest building block of matter

John Dalton’s Model

“An atom is comprised of many tiny spherical particles which are indestructible and indivisible.” - Deduced in 1802

Elements

Atoms containing only one type of atom

Compound

Atoms containing many different elements

Protons, Neutrons, Electrons

Positively charged, neutrally charged and negatively charged particles within an atom.

Size of electron

8000 times smaller than a a proton

Electrostatic attraction

Positive electrons attracting negative electrons

Nucleons

Protons & Neutrons

Flame Test (Purpose)

To distinguish metallic elements when put in a flame.

Order of Energy Emitted In Terms Of Color

Violet, blue, green, yellow, orange, red

Bohr Model

“Electrons move around in fixed orbits around the nucleus, electrons have fixed energy levels, electrons cannot exist between two energy levels, larger orbits correspond with larger energy levels.” - Deduced in 1913, by Niels Bohr

Electron Shells

Energy levels where electrons are grouped in.

Effects Of Heating An Element

Electron jumps to a higher energy state, and the energy is then emitted by light or heat, as the electron returns back to its original state.

Ground State

Lowest energy level for an atom

Excited Energy State

Electron when “excited”.

Formula for no. of electrons per shell

2n²

Problems With Bohr Model

• Unable to explain why shells can only hold 2n² electrons

• Can’t predict emission spectra for atoms that have more than one electron

• Can’t explain why 4th shell gain 2 electrons before 3rd shell fills up.

New Concept In the Schrodinger Model Vs. Bohr Model

Concept of “orbitals” and “subshells'“

Subshells & Oribitals

Subshells: Shells that contain separate energy levels (s,p,d,f)

Orbitals: Smaller components within the subshell.

Electrons To Be Held In An Orbital (Max)

2

Orbitals present in s,p,d,f

1,3,5,7 (possible electrons being: 2,6,10,14)

Energy order of subshells

1s,2s,2p,3s,3p,4s,3d

Exceptions For Rule Of Orbitals (Elements)

Chromium, Copper

Electron Configuration for Chromium & Copper

Chromium: 1s²,2s²,2p^6,3s²,3p^6,3d^5,4s1

Copper: 1s², 2s², 2p^6, 3s², 3p^6, 3d^10, 4s^1

Reason For Copper & Chromium Electron Arrangement

It is slightly more stable in this configuration.

Main Group Elements In Periodic Table

Group 1,2 13-18 (Excluding Transition Metals)

Reason For Elements Having Similar Properties In Same Group

Have same no. of electrons

Properties Of Alkali Metals

Soft and reactive with oxygen and water.

Properties Of Noble Gases

Low reactive

Have a full Outer shell

Information stated by periodic arrangement

Number of shells, and is in the same column as other elements with same no. of shells

Blocks Within Periodic Table

These blocks indicate which elements have their valence electron in what subhell. Eg. “s” blocks elements have their valence electrons in the ‘s’ block.

Electronegativity

Ability for an atom to attract an electron (higher rating corresponds to higher ability)

First Ionization Energy

Energy required to remove one electron from the atom.

Person who observed periodic trends

Dimitrivi Mendeleev - 1869

Reason For Electrostatic Attraction Decreasing Down A Group

More shells, less electrostatic attraction retaining them together.

Effective Nuclear Charge/Core Charge

“Attraction felt by the valence electrons from the nucleus.”

Found by: no. of valence electrons

Core Charge Across A Period

Increases - due to more valence electrons

Electronegativity - down the group and across the period

Decreases Down Group: Electrons become more distant from nucleus despite core charge remaining the same.

Increases Across Period: Number of valence electrons increase, thus increasing core charge - increasing electrongeativity.

Atomic Radius - down group and across period

Increases down the group: More shells thus higher atomic radius.

Decreases down the period: More valence electrons - more electrostatic force - less atomic radius.

First Ionization Energy - Across Period, Down Group

Across Period - Increases: More electrostatic force holding together electrons due to higher core charge.

Down Group - Decreases: Less electrostatic force, as more shells are being formed and less electrostatic attraction.

Most Reactive Metals

Group 1 - Alkali Metals

Properties For Metals Down A Group

More Reactive - More Shells, and less core charge and electrostatic attraction. Electrons can easily be lost without a lot of ionization energy - thus more reactive.

Non-metal electron gaining properties

Across Period: More reactive

Down Group: Less reactive as they gain electrons unlike metals.

Most Reactive Metal

Francium

Most Reactive Non-Metal

Fluorine

Thermal Conductivity

The ability to transfer heat

Non-metals reactivity

Decreases down a group - harder to attract electrons down a group

Increase across the period