Name:_______________________ Unit 3: The Atom1) Atomic Structure (The Nucleus) (HW: p. 19, 20)Essential Question: What is everything made up of?

Atoms are the smallest pieces an element can be broken into and still retain the properties of that element. It comes from the Greek word atomos, meaning “indivisible” (unbreakable).

Atoms are so tiny that they can not be seen directly. They can be detected through X-ray crystallography or atomic force microscopes, but only indirectly.

It takes 602 000 000 000 000 000 000 000 atoms of hydrogen to weigh 1 gram (the mass of a small paper clip).

Atoms are made up of the following particles:

A) Nucleons (Particles in the Nucleus)

1) Protons: have a mass of 1 atomic mass unit (1.66 X 10-24 grams) and a charge of +1. They are found in the nucleus of the atom, and the number of protons in the atom is the atomic number, which identifies what element the atom is. Oxygen (O) has an atomic number of 8, which means there are 8 protons in the nucleus. Since protons are the only particle in the nucleus to have a charge, the charge of the nucleus is + (# of protons). Since oxygen has 8 protons in the nucleus, oxygen has a nuclear charge of +8. We will use nuclear charge down the road for the purposes of explaining why it is easier to do nuclear fusion with smaller nuclei and why atoms have the sizes they do.

2) Neutrons: have a mass of 1 atomic mass unit, and no charge. They are found in the nucleus of the atom, and the number of neutrons added to the number of protons gives you the mass number of the atom. The number of neutrons does not affect the identity of the element. Oxygen’s most common form has a mass number of 16. Since there are 8 protons in the nucleus of oxygen, this means there must also be 8 neutrons to give a combined mass of 16. The number of protons and neutrons does NOT have to be equal. In addition, atoms of any given element can have differing numbers of neutrons. Atoms of the same element with different numbers of neutrons in their nuclei are called ISOTOPES of one another. The most common isotope is the one who’s mass equals the average atomic mass given on the periodic table rounded to the nearest whole number. Since O has a given average mass of 15.9994, the most common isotope of O is O-16, or Oxygen with a mass number of 16. See the diagram below.

B) Particles Outside The Nucleus

3) Electrons: have a mass of 1/1836 amu (9.11×10−28 grams) and a charge of -1. They are found orbiting the nucleus in energy levels. Atoms gain, lose or share electrons when they form chemical bonds. If electrons are gained and lost, an ionic bond is formed. If electrons are shared, a covalent bond is formed. The number of electrons in the atom equals the number of protons. Atoms are neutrally charged, so the + charged protons and the – charged electrons must be equal in number to give a neutral charge. Oxygen has 8 protons in its nucleus, so there must be 8 electrons zipping around outside the nucleus in energy levels. In this unit, the only thing you need to worry about is how to find out how many electrons an atom has, and what the charge and mass of an electron are. Later in the course, you will see just how important electrons are to all of chemistry. They are the part of the atom responsible for all chemical bonding. If it weren’t for electrons, there would be no compounds. H would not bond to O, and water would not exist!

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1) How many protons does an atom of iron (Fe) have? With an atomic number of 26, the atom has 26 protons.

2) What is the nuclear charge of an atom of Fe? Since Fe has 26 protons, the nuclear charge is +26.

3) How many atomic mass units (amu) do the protons in an atom of Fe weigh? Since each proton has a mass of 1 amu, the mass of 26 protons is 26 amu.

4) How many electrons does Fe have around its nucleus? The number of electrons in the atom must equal the number of protons. Since Fe contains 26

protons, it must also contain 26 electrons. If you look at the electron configuration 2-8-14-2 and add up all of those electrons, you will see it adds up to 26 electrons.

5) What is the most common isotope of Fe? Take the average atomic mass of all the isotopes of Fe (55.847) and round it to the nearest whole number (56). That is the mass number of the most common isotope of iron, Fe-56.

6) How many neutrons are in the nucleus of the most common isotope of Fe? To find the number of neutrons, take the mass number (which is the number of particles in the nucleus, protons and neutrons combined) and subtract out the atomic number (the number of protons). There are 56 particles in the nucleus of the most common isotope (Fe-56), of which 26 are protons (from the atomic number of 26). 56 – 26 = 30 neutrons in the nucleus.

OK, what about these isotopes? Why does Fe have an AVERAGE atomic mass of 55.847? What are the other isotopes of iron?

There are four common naturally occurring isotopes of iron:

Mass of Isotope(amu)

Notation 1

Symbol – mass #

Notation 2

Mass #Symbol

# protons(atomic

#)

# neutrons(mass # - atomic

#)

# electrons(atomic #)

% Abundance in Nature

54 Fe-54 54Fe 26 54-26 = 28 26 5.845%56 Fe-56 56Fe 26 56-26 = 30 26 91.754%57 Fe-57 57Fe 26 57-26 = 31 26 2.119%58 Fe-58 58Fe 26 58-26 = 32 26 0.282%

As you can see, the most common (abundant) isotope of iron really is Fe-56! It makes up more than 90% of all iron atoms.

But that average…looks kind of suspicious! If you average 54, 56, 57 and 58, you get 56.25, not 55.847. Apparently the average mass is not a straight average, the type you are used to calculating! See the next page for the sordid details!

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Calculating Weight-Average Atomic Mass

The atomic masses given on the periodic table are WEIGHT-AVERAGED masses. This is calculated using both the masses of each isotope and their percent abundances in nature. For the purposes of simplicity, we will round weight-average mass to the THOUSANDTHS place.

To find the weight-average mass of an element given the mass of each isotope and each isotopes percent abundance:

WAM = (massisotope 1 X %/100) + (massisotope 2 X %/100) + (massisotope 3 X %/100) + …So, for the four isotopes of iron:

Weight average mass of Fe is:

WAM = (54 X 5.845/100) + (56 X 91.754/100) + (57 X 2.119/100) + (58 X 0.282/100) WAM = (3.156) + (51.382) + (1.208) + (0.164)

WAM = 55.910 amu

Another way to look at this is as follows: take 5.845% of 54 (which is 3.156), then take 91.754% of 56 (which is 51.382), then take 2.119% of 57 (which is 1.208) and then finally 0.282% of 58 (which is 0.164). Take the four numbers you get as a result and add them together, and you have the weight-average mass!

Note that this is still not exactly the same as the listed weight-average mass on the Periodic Table. The isotope information you used to solve this problem came from the National Nuclear Data Center at Brookhaven National Laboratory. The weight-average mass given on the Periodic Table may include the other isotopes of iron, which are all radioactive and make up a very tiny percentage of iron’s mass. The weight-average mass is based on the abundance of the naturally occurring isotopes of that element.

Also, protons and neutrons do not weigh exactly 1 amu. Neutrons weigh a tiny fraction more than protons do. An atomic mass unit is actually an average mass, found by taking the mass of a C-12 nucleus and dividing it by 12.

Is how it is shown on the Periodic Table’s key.

Here’s another example problem:Boron (B) has two naturally-occurring isotopes. B-10 has a mass of 10.0 amu and makes up 19.80% of all B atoms. B-11 has a mass of 11.0 amu and makes up 80.20% of all B atoms. What is the weight-average mass of boron?

WAM = (10.0 X 19.80/100) + (11.0 X 80.20/100) = 10.802 amu

Basically, you are taking 19.80% of 10.0, then taking 80.20% of 11.0 and adding the two numbers together.

Mass of Isotope(amu)

Notation 1

Symbol – mass #

% Abundance in Nature

54 Fe-54 5.845%56 Fe-56 91.754%57 Fe-57 2.119%58 Fe-58 0.282%

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Where are electrons found in the electron cloud?

I. The Bohr Atom1. Bohr was the first to propose that the electrons were located in energy levels. A lower case “n” is used to denote these principle energy levels (also called principle quantum numbers). The principle energy levels are numbered, so that the level closest to the nucleus is labeled n = 1. The next level is labeled n = 2 and so forth. Each principle energy level had a certain energy value associated with the level. The closer the level was to the nucleus, the lower the energy of the level. The further away from the nucleus, the higher the energy is of that level. As long as the electrons were in these levels, the electrons do not give off energy. The dark circle below represents the nucleus. The rings around the nucleus represent the principle energy levels. Number the principle energy levels starting with the one closest to the nucleus: n = 1, n = 2, n = 3 etc.

2.

Electron Configuration and the Periodic Table

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Each principle energy level can only hold so many electrons before the level is full. A quick and easy way to determine the maximum number of electrons (max e-) that a principle energy level can hold is given by the following: max e- = 2 n2. First square the principle energy level number (n) then multiply by 2.

Energy Level (n)

Maximum number of electrons (max e- = 2 n2)

123456

Electrons are arranged around the nucleus by filling up the first principle energy level (n=1), then the second energy level, etc. This is the electron configuration given on your periodic table. The number of electrons are listed for each level with a dash between levels: for oxygen (O) which has a total of 8 electrons, the configuration is 2–6 (2 electrons are located in the first principle energy level and 6 electrons are located in the second principle energy level. Look up the electron configuration on the periodic table for the element given and fill in the chart. Ca is done as an example.

Element n = 1 n = 2 n = 3 n = 4Ca 2 8 8 2NaFBAlCH

3. Completely Filled vs. Occupied Principle Energy LevelsOccupied means that there is at least one electron in the Principle Energy

Levels6

Li: 2 – 1 has 2 occupied Principle Energy LevelsCompletely Filled means that each level has its maximum number of

electrons which can be determined by the 2n2 rule.Li: 2 – 1 has only 1 Completely Filled Principle Energy Level

To help you review the 2n2 rule complete the following chart

PEL (n) 1 2 3 4 5 6 7Max e-

For the Following:a) Copy the electron configuration from the Periodic Tableb) Determine the number of Occupied Principle Energy Levels (PEL)c) Determine the number of Completely Filled Principle Energy Levels

Element Electron Configuation

# Occupied PEL

# Completely Filled PEL

C

Na

O

Cl

He

F

Ne

Si

Zn

Au

4. Drawing Bohr Diagrams of Atoms: 1) A circle is used for the nucleus- the # protons (# p or +) and the # of neutrons (#n) are placed in the circle.2) A ring is drawn around the nucleus for each energy level. 3) The electrons for each energy level are placed in pairs symmetrically around the nucleus

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For F: atomic # = _____________atomic mass = _____________electron configuration: ______________# p = _________ # n =____________

For Al : atomic # = _____________atomic mass = _____________electron configuration: ______________# p = _________ # n =____________

Going Backwards: Determining the identity of an element from the Bohr diagram:

# p = _____________ # n =______________atomic # = _____________atomic mass = # p + # n = ________________ electron configuration: ____________________________________Isotopic Notation:

# p = _____________ # n =______________atomic # = _____________atomic mass = # p + # n = ________________ electron configuration: ____________________________________Isotopic Notation:

II. Introduction to LightVisible Light (energy we see with): part of the Electromagnetic Spectrum

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1. Two theories to explain light’s behavior:

Waves Particles of Packets of Energy

There was evidence for both models so the two theories were put together!!

Light: QUANTUM THEORY OF LIGHTa) packets or bundles of energy called _________________ or ______________b) travel in wave-like fashionc) produced when electrons drop from ______________ energy levels to _______

energy levels (the greater the drop, the greater the energy the light has)

2. Properties of Light

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Wavelength () - ____________________________________________Frequency (F) - ______________________________________________

(units: Cycles / second OR Hertz)

Energy (E) - __________________________________Speed (velocity) – same for all electromagnetic radiation _______________

Relationships: Frequency and Energy: Type _________________F ______, E ________ or F ______, E ________

Frequency and Wavelength: Type _________________F ______, ________ or F ______, ________

Wavelength and Energy: Type _________________ ______, E ________ or ______, E ________

3. Bright Line Spectra and Continuous Spectrum A. Bright Line Spectra and the Bohr Atom

An electron must absorb energy before it can give off colors we see in the bright line spectra. When energy is added, the electron moves to a higher energy level. The potential energy of the electron increases. This is an unstable situation. In order for the electron to return to a lower and more stable energy level, the added energy must be given off. When the electrons return to the lower energy levels this decreases the PE because the added energy is given off and the colors of the bright line spectra are seen. Moving electrons to different energy levels requires different amounts of energy. These different amounts of energy produce the different colors.

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Movement of an electron between the same 2 energy levels in DIFFERENT elements will produce different colors. The energy between the energy levels depends on the number of protons and the number of electrons that each element has.

BRIGHT LINE SPECTRA are produced when “electrons in the EXCITED STATE” fall back to lower energy levels of the GROUND STATE. Unlike the continuous spectrum of sunlight, only certain colors will be present in the BRIGHT LINE SPECTRA. The BRIGHT LINE SPECTRUM is like a “fingerprint” of the element that produced the spectrum. Like a fingerprint, the BRIGHT LINE SPECTRA can be used to identify the element. When viewed with a spectroscope, the individual bands of colors in the BRIGHT LINE SPECTRUM can be seen and the wavelength of each band determined.

1. Below are the BRIGHT LINE SPECTRA of three elements. From the position of the lines determine which element is the unknown. (HINT: Match up the lines present in the unknown with the three known elements.) Unknown element = ____________

Element X

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Element Y

Element Z

Unknown

2. Which of the two elements above are present in the BRIGHT LINE SPECTRUM given below? (HINT: Match up the lines present with the three known elements. Only two patterns should match perfectly.) ____________ and _____________

B. The Rainbow: A Continuous Spectrum

Long Short Low F High FLow E High E

LONG STEM RED ROSES: All “L’s” go together with RED

Continuous Spectrum

when radiation from the sunlight passes through a prism, a rainbow – a spectrum of colors – is seen the colors are not separated from one another but blend together due to the overlap of the line spectra of the 67 different elements in the sun

lightbulb

C. Bright Line Spectrum

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ROYGBIV

R O Y G B I V

when radiation from an excited atom (element) passes through a prism, the radiation is separated into various wavelengths and colors

Colors are not blended – spectrum is discontinuous – and you observe lines of color at different locations

Flame

5. Ground and Excited StatesA) The lowest possible energy state that an electron can occupy is called the __________ ____________. This is a very __________ condition. The principle energy levels, which are occupied match those predicted by the electron configuration on the periodic table. When electrons gain energy, the electrons move to higher principle energy levels then they would normally occupy. This unstable situation is called the _____________ _______________. The electrons will release the absorbed energy, often seen as the bright line spectrum of the element, and fall back to the ground state.

A) How to tell when energy will be absorbed or releasedThe Principle Energy Level (n) changes:

If the number of the principle energy level (n) goes up, then energy is _____________ or ______________ n = 1 to n = 3 OR n = 3 to n = 4

If the number of the principle energy level (n) goes down, then energy is _____________ or ______________ n = 2 to n = 1 OR n = 5 to n = 3

If the energy is emitted, then ____________________ (colors) are seen.Determine if energy is added/absorbed (+E) or released/emitted (-E) for the following transitions ; circle the

1) n = 1 to n = 2 ______________ 6) n = 1 to n = 5 ______________

2) n = 4 to n = 3 ______________ 7) n = 4 to n = 2 ______________

3) n = 2 to n = 1 ______________ 8) n = 2 to n = 3 ______________

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ROYGBIV

B) How do you tell the excited and ground state apart from the electron configuration??Ground State: Matched the predicted electron configuration found on the periodic table. In other words, it follows the order given

Ground State for Oxygen (O) on PT= 2 – 6 (8 total electrons)

Possible Excited State for Oxygen = 1 – 7 (still 8 total electrons)The first energy level is not filled before moving into the second energy level.

The KEY here is that the configuration

does not MATCH the one on the PT. Another possible excited state for oxygen: 2 – 5 – 1 (still 8 total electrons)For the following elements, fill in the chart and determine if the electron configuration is in the GROUND STATE (GS) or EXCITED STATE (ES).

III. UpDaTiNg ThE BoHR MoDeL: QuAnTuM oR WaVe MeChAnIcaL MoDeL

Element Symbol

Electron Configuration

Principle Energy Levels

Ground or excited state?

C 2-4F 1-8Cl 2-8-7B 2-1-1-1Na 2-7-2O 2-6S 1-8-7Zn 2-8-18-2Br 1-7-17-8-1-

1Br 2-8-18-7

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The Bohr Model was very good at explaining many things about atoms and electrons and how they behaved. Just like your wardrobe needs updating, so did the model of the atom. A revolution in physics occurred in the early 1900’s when experiments showed that matter, just like light energy, could have a dual nature…it can act as a particle or a wave.

1. The Wave Mechanical Model A. Still has a dense positive nucleus where protons and neutrons are locatedB. Using complicated math, this math showed that electrons were not

moving in definite fixed orbits like planets but had distinct amounts of energy.

C. Four “Quantum Numbers” are used to describe the location of the electron.

2. Four “Quantum Numbers”1) Principle Energy Levels (PEL) are also called the PRINCIPLE QUANTUM NUMBER.) Same as the Bohr Energy levels (n = 1, 2, 3, 4, 5, 6, 7)

2) Sublevels: PEL get divided into sublevels. Like Different room types- gives the 3-D shape where electrons are found.

These sublevels are designated: s, p ,d and f The number of sublevels is determined by the number of the

princple energy level (n): n = # sublevels Within the same principle energy level: the sublevels have

different energies: s p d f

lowest energy highest energy

3) Orbitals are “regions” or “areas” around the nucleus The sublevels are divided into orbitals (rooms) where electrons are found. Each orbital can hold 2 electrons.

The number of orbitals depends on the sublevel type: The number of electrons in a sublevel depends on the

number of orbitalsSublevel # orbitals x 2 = # of electrons in sublevel

s x 2 =p x 2 =d x 2 =f x 2 =

An orbital is defined as a region in which an electron with a particular amount of energy is most likely to be found.This volume outside the nucleus are often described as an Electron

Cloud.15

4) SPIN: or Within an orbital, the two electrons spin in opposite directions to overcome the repulsion they feel for each other (remember- like charges repel.)

Principle Energy Level

Sublevel(s) Present

# Orbitals Present

# electrons in Sublevel

Total # of e per PEL

1

2

3

16

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Remember the tricks for number of sublevels, orbitals and total number of electronsFor Principle Energy Level (PEL) n

n = _________________________ & n = ___________________________

n2 = _____________________ & 2 n2 = ________________________

2. Quantum Atom Electron Configuration from the Periodic Table There are two ways to show the Quantum Atom Electron Configuration. One way shows the number of electrons in the Principle Quantum Level (PEL) and Sublevel type (letter). The other way, known as BOX DIAGRAMS, show all four quantum numbers by using boxes for orbitals and arrows for electrons to show the opposite spin of the electrons. Using the electron configuration from the periodic table, we will do both types.For Hydrogen: EC form PT is 1

Quantum Atom Box Diagram 1s1 # of electrons 1s ORBITAL

PEL SUBLEVEL PEL SUBLEVELARROW SHOW SPIN

RULES TO REMEMBER WHEN DRAWING BOX DIAGRAMS; * S has one orbital so draw 1 box

* P has 3 orbitals so draw 3 boxes, * Each “P” box in a sublevel must have an electron before pairing up because each orbital has the same energy as the other orbitals (Boxes)* Electrons spin in opposite directions!

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Element

Elec. Config. From P T

Quantum Atom E. C.

Box Diagrams

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H

He

Li

Be

Element

Elec. Config. From P T

Quantum Atom E. C.

Box Diagrams

B

C

N

O

F

Ne

Na

Mg

Al

Si

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P

Element

Elec. Config. From P T

Quantum Atom E. C.

Box Diagrams

S

Cl

Ar

K

Ca

After argon (Ar) the energy levels begin to overlap and things get complicated. We will not be doing elements above Calcium. However….

How many orbitals are in the d sublevel? _________ f sublevel? _____________

How many boxes would you draw for the d sublevel? _________ f sublevel? __________

IV. The Kernel, Valence Electrons and Lewis (Electron) Dot Diagrams

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A very powerful tool for showing how different elements bond together, called Lewis (Electron) Dot Diagrams, was developed by a chemist named Lewis. Many times in chemical bonding, only the electrons on the outside or in the highest energy level, called the VALENCE SHELL, actually become involved in bonding to form compounds. We are going to look at the Lewis (Electron) Dot Diagrams and elements in this next section.

Valence Shell: Outermost energy level of an element (highest number)Oxygen: O 2-6 the 2nd principle energy level is the valence shell

Valence Electrons (VE): electrons located in the outermost energy level- these are the electrons which are the furthest to the right in the electron

configuration found on the PT- Oxygen: O 2-6 the 6 electrons are the valence electrons- In a dot diagram, the VE are represented by dots:

Kernel: This is the nucleus and all the other electrons located in the inner energy levels

In a dot diagram, the chemical symbol represents the kernel: OThe valence shell always contains only 4 orbitals. These orbitals are represented by the four sides around the symbol. The top side represents s orbital which has less energy than the other three p orbitals. The other three p orbitals have the same amount of energy. First 2 electrons are always placed first in this top orbital. Going around clockwise, one electron is placed on each remaining side until you need to pair the electrons up. (You don’t draw the OVALs!)

s Orbitallower energy than others

O O

the same energy the same energy

p orbitals p oribitalsSymbol Li Be B C N O F NeElectron

Configuration

Dot Diagram Li Be B C N O F Ne

Guided Practice: look for a pattern!Symbol Electron Configuration (Circle Number of VE Dot Diagram

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the Valence Electrons (VE)

S

Se

Na

K

Si

Ge

Cl

Br

Ar

Kr

What was the pattern for electron dot diagrams?______________________________________________________________________________________________________________________________________

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