Unit 2 - Atomic Theory

UNIT 2 - ATOMIC THEORY

VOCABULARY:

Allotrope Electron Configuration Nuclear Charge Anion Element Nucleons Atom Excited state Nucleus Atomic Mass Ground state Orbital Atomic Mass unit (a.m.u.) Ion Proton Atomic number Isotope Quantum Theory Bohr model Kernel electron(s) Valence electron(s) Cation Lewis Dot Diagram Wave-mechanical model Compound Mass number Electron Neutron

OBJECTIVES: Upon completion of the unit you will be able to do the following: • Understand that the modern model of the atom has evolved over a long period of time through the work of many scientists • Discuss the evolution of the atomic model • Relate experimental evidence to models of the atom • Identify the subatomic particles of an atom (electron, proton, and neutron) • Know the properties (mass, location, and charge) of subatomic particles • Determine the number of protons, electrons, and neutrons in a neutral atom and an ion • Calculate the mass number and average atomic weight of an atom • Differentiate between an anion and a cation • Identify what element the amu unit is derived from • Define the term orbital • Distinguish between ground and excited state • Identify and define isotopes • Write electron configurations • Generate Bohr diagrams • Differentiate between kernel and valence electrons • Draw Lewis Dot Diagrams for an element or an ion 1

THE EVOLUTION OF THE ATOMIC MODEL

Atom = Democritus = Dalton (1803) o Known as the ______of the atomic theory o Dalton invented the word ______as the basic unit of matter which were considered to be ______o Dalton also claimed that all atoms of a given element are ______o He also discovered that atoms of different elements have different ______o Found that combining atoms of different elements formed ______o Theory referred to as the ______theory (it looked like a simple sphere) *What does this name tell you about Dalton’s atom?

J.J. Thomson (1897) o While using a ______he discovered that the ray was deflected (due to a magnetic/electrical field) o From this discovery he concluded that atoms contain small negatively charges particles called ______o Theory famously referred to as the ______model because he visualized the ______being ______within the structure of the atom (just like raisin bread) o The ______of the rest of the atom (besides the electrons) was thought to be ______and ______

* Rutherford (1909) * o Experiment called the ______where he ______a thin piece of ______with a ______o Often referred to as the ______model o Most alpha particles went ______& some were ______o Two conclusions were therefore made: 1) most of the atom is ______2) atoms have a ______, ______called the ______

* Neils Bohr (1913) *

o Proposed that the atom consists of a dense nucleus with ______found in ______o He therefore stated that each electron orbiting the nucleus must possess a ______to keep it in place within its orbital o Known as the ______model (looks much like our solar system)

Wave-Mechanical/Cloud Model (Modern Present day model) o Developed after the famous discovery that energy is made up of BOTH ______& ______o Still the same dense positively charged ______o Electrons now have distinct amounts of energy and move in areas called ______or ______o ______have contributed to this theory o Different from the Bohr diagram—now the location of the electron is based on ______within the orbital *3-dimensional model 3

Atomic Model Conclusion(s) Date: Scientist: Name of Model:

Date: Scientist: Name of Model:

Atomic Model Conclusion(s) Date: Scientist: Name of Model:

Date: Scientist: Name of Model:

VOCABULARY (of the Periodic Table)

SUBATOMIC PARTICLES Subatomic Charge Relative Mass Location Symbol How to Calculate Particle

Proton

Neutron

Electron

Let’s Practice Calculating Subatomic Particles in Different Atoms:

Symbol # # # Atomic Mass Nuclear Protons Neutrons Electrons Number Number Symbol

35 Cl 17

15 16

C-14 8

Mass number =

Nuclear Charge =

Nucleons =

ATOMS (neutral) VS. IONS (charged)

Vocabulary Term Definition Example/Diagram

Neutral Atom

Ion Anion

Cation

ISOTOPE =

Example: Isotopes of Carbon (C-12, C-13, & C-14)

12 13 14 C C C 6 6 6

U-238 U-240

So why does Carbon have a mass of 12.011 on the Reference Table? This is carbon’s ATOMIC MASS which is the average weighted mass of all naturally occurring isotopes of carbon (there exist three different isotopes of carbon in the atmosphere – as seen just above)

Calculating Atomic Mass (for any element):

Atomic Mass = the weighted average of an element’s naturally occurring isotopes

(abundance in decimal form) (mass of isotope 1) (abundance in decimal form) (mass of isotope 2) + (abundance in decimal form) (mass of isotope 3)

Example 1

12 C = 98.89% of carbon in the atmosphere 13 C = 1.11% of carbon in the atmosphere

Step 1: Multiply the mass of each separate isotope by its percent abundance (in decimal form!!!!!!!)

Step 2: Add up the products of all the calculated isotopes from step 1.

Example 2: The element Boron occurs in nature as two isotopes.

Isotope mass percent abundance Boron 10.0130 amu 19.9% Boron 11.0093 80.1%

Example 3: Isotope of Hydrogen Percent Abundance 1 H Protium 99.0% 1 2 H Deuterium 0.6% 1 3 H Tritium 0.4% 1

ATOMIC MASS VS. MASS NUMBER

ELECTRON CONFIGURATIONS = a dashed chain of numbers found in the ______of an element box (see below); tells us the number of ______as well as the number of ______in each level (tells us how the electrons are arranged around the nucleus)

**All electron configurations on the Periodic Table are NEUTRAL (p=e)

Substance Electron Configuration Magnesium Mg +2 Bromine Br -1 Barium *Lead

* shortcut allows you to cut out the first two orbitals to shorten the “address”

Valence Electrons: electrons found in the ______shell or orbital; the ______number in the electron configuration

Kernel Electrons: ______electrons (all non-valence electrons)

Sulfur # valence e- ____ Nitrogen # valence e- ____ # kernel e- ____ # kernel e- ____

Principle Energy Level (n) = electron energy levels that contain a certain number of ______; each sublevel contains one a set number of ______

Maximum # of electrons in an energy level = ___ where n = ______# (or period #)

Principle Energy Level (n) Maximum number of electrons ( ) 1 2 3 4

Sample question: What is the maximum number of electrons that can occupy the 3 rd principal energy level in any atom? ______

BOHR DIAGRAMS (one method for expressing electron location in an atom or ion); ______MUST be drawn ______

1. Look up electron configuration of element at hand on Periodic Table (if you are working with an ion, add/subtract the proper amount of electrons from outer shell(s) of configuration)

Example: Carbon is ______

2. Draw nucleus (with a circle) and notate correct amount of protons and neutrons inside

3. Using rings or shells (these are your orbitals), place the proper amount of electrons in their appropriate orbital(s) – there should be as many rings/shells as dashed number in electron configuration

Carbon Fluorine Beryllium Al

Li I Na + S-2

LEWIS (ELECTRON) DOT DIAGRAMS - a shorter way of expressing electron location; an ______

• Only illustrates ______!

1. Write the element’s symbol

2. Retrieve electron configuration from Periodic Table. The last number in the configuration is the ______.

3. Arrange the valence electrons (DOTS) around the symbol using the following rules: • Only two electrons maximum per side of the symbol (therefore no more than 8 total surrounding symbol – 8 is great!) • Always “pair” the first two • If you have more than 2 valence electrons, deal them one at a time to the other three sides until you run out 1 2 X 4 7 4. If you are working with an ion you must adjust the valence electrons (add or subtract electrons) in the configuration before constructing your Dot Diagram – be sure to draw your final diagram with the initial charge on the ion

Draw LEWIS ELECTRON-DOT DIAGRAMS for the following:

Argon Phosphorus Carbon Beryllium Oxygen

Aluminum Sodium Bromine Na +1 O-2

15 Ground State vs. Excited State

*Notice that one electron from the 2nd orbital has moved to the 3rd orbital

Ground State = electrons in ______possible (configuration found on ______); electrons ______as physically possible ground state electron configuration for Li is _____

Excited State = electrons are ______(____ configuration ____ found on PT) excited state electron configuration for Li could be ______

Distinguish between ground state and excited state electron configurations below:

Bohr Electron Configuration Ground (G) or Excited (E) state? 2-1 2-0-1 1-1-1 2-7-3 2-8-2 2-8-8-2 2-8-17-6 2-8-18-8 2-6-18-1 2-5-18-32

The greater the distance from the nucleus, the greater the energy of the electron

Ground excited energy is absorbed Ground  Excited energy is released (in the form of light energy)

When atoms ______their ______will shift to a ______energy level or ______. • This is a very ______condition so the electrons will ______state energy level (or ______) • When they ______from the excited state to the ground state they release energy in the form of ______. Different elements produce different colors of light or ______. These spectra are ______for each element (just like a human fingerprint is unique to each person). We therefore use spectra lines to ______which element we have because each element gives off a characteristic bright line spectra.

What to Study for the Atomic Exam (Unit 2)

Structure of the atom (role/nature of the nucleus and electrons etc.) Gold Foil Experiment (know what was observed and the conclusions that were drawn from these 2 observations) Orbit vs. orbital (what is the difference?) Location, mass, charge of protons, neutrons, and electrons in an atom (see chart in notes) Isotopes (same element; different number of neutrons or masses ex: C-12 & C-14) – be able to identify them Nucleons (the subatomic particles found in the nucleus) Nuclear Charge (charge inside the nucleus; always positive and directly dependent on # protons) Chemical notations: where do we find the atomic #, mass #, atomic mass etc. Determining the number of protons, neutrons, and electrons (pne) in an atom or ion (use mass # or atomic mass rounded to whole number) – remember that the # protons ALWAYS tell us the symbol/element we have Atomic mass – weighted average mass of an element’s naturally occurring isotopes (ex: C is 12.011) Ground state vs. excited state and the energy absorbed/released during electron movement (remember: light is emitted or seen when an electron jumps from an excited state back down to ground state or a lower energy level) Principle energy levels (orbitals) and the maximum # electrons found in each – use 2n 2 formula Calculating average atomic mass given isotope abundances and masses The nucleus makes up pretty much all the mass in an atom – therefore (mass #) = (#protons) + (#neutrons) Who performed the cathode ray experiment? What subatomic particle was discovered in this experiment? Atom vs. Ion (what is the difference?) Electron Configuration – found below the atomic number on PTable (know how to write them for an atom or ion) Drawing Bohr models – show all the electrons within an atom/ion and their location (construct nucleus & use shells or circles to place electrons) Lewis Dot Diagrams – write element symbol and distribute dots around (draw only the valence electrons)