Chem A UNIT 2 REVIEW - Objectives

Lesson: Early Chemistry and Atomic Structure

Describe the changes in the atomic model over time. Explain how different observations and experiments led to changes in the atomic model.

Lesson: Properties of the Atom

Describe the properties and locations of protons, neutrons, and electrons in an atom. Explain what the mass number and atomic number identify about an atom.

Use mass and atomic number to determine the number of each subatomic particle in an atom.

Lesson: Isotopes and Weighted Averages

Identify the differences between isotopes of an element. Calculate the weighted average mass of an element when given the percent abundance and mass of its

isotopes.

Lesson: The Mole and Conversions

Describe the mole and its use in chemistry calculations. Use the mole, Avogadro’s number, and molar mass in dimensional analysis calculations.

Lesson: Electron Arrangement and Electromagnetic Radiation

Describe observations that led to the Bohr and quantum models of the atom. Explain how line spectra are used to identify elements and what they indicate about atoms.

Explain properties of waves and how these properties apply to the quantum model.

Describe your observations, calculations, and conclusions from the virtual lab.

Lesson: Electron Configurations

Represent electron arrangements using electron configuration, orbital notation, shorthand notation, and Lewis dot notation.

Describe the quantum model of an atom in terms of energy levels, sublevels, and orbitals.

Early Chemistry and Atomic Structure: History of the Atom—Text VersionRead with a purpose …Describe at least two changes that occurred to the atomic model over time, along with the experimental evidence that made these changes necessary.As you explore the historical development of the atomic model, pay particular attention to how theories and models are proposed, developed, and modified to account for new experimental observations and results.Early Ideas of the Atom:As early as 400 B.C., some Greek philosophers had the idea that matter could be divided into a basic particle of matter. The Greek thinker Democritus called these particles "atoms," from the Greek word meaning "cannot be cut into smaller pieces." Dalton's Atomic Theory:John Dalton, proposed an atomic theory built upon experimental results. Dalton's atomic theory still serves as a foundation for the current atomic model, even though exceptions to his theory are now known.  His atomic theory can be summarized by the following statements.

1. All matter is composed of extremely small particles called atoms. Scientists still agree that all matter, from the air you breathe to the computer you are using, is composed of atoms.

2. Atoms of a given element are identical in size, mass and other properties. Scientists now know that all atoms of a given element have the same number of protons. However, you will learn in a later lesson that atoms of the same element can have different masses because of a different number of neutrons.

3. Atoms cannot be subdivided, created, or destroyed in chemical reactions. Scientists now know that atoms are divisible and made up of smaller, subatomic particles. However, the law of conservation of mass still states that matter cannot be created or destroyed in chemical reactions. Atoms of different elements can combine in simple, whole-number ratios to form chemical compounds

4. Atoms of different elements can combine in simple, whole-number ratios to form chemical compounds. Scientists still agree that atoms combine in whole-number ratio to form compounds like CO2 and H2O.

5. In chemical reactions, atoms are combined, separated, or rearranged. Scientists still agree that bonds are broken and formed as atoms are rearranged in chemical reactions.

The first model of the atom, proposed by Dalton, was a solid sphere. Dalton believed that atoms were the smallest form of matter. Although scientists now know of exceptions to Dalton's theory, it has not been discardedSir William Crookes:In one of his experimental investigations with cathode ray tubes, Crookes placed an object within the path of the cathode ray. Because a shadow appeared on the positive electrode, Crookes concluded that something (either light or small particles) was traveling from the negative to the positive electrode. Crookes hypothesized that the glow was made up of small particles of matter, Thomson and the Electron:Sir J.J. Thomson conducted more experiments using cathode ray tubes by placing a negatively or positively charged rod near the cathode ray tube. The beam is affected by the electrically charged rods. Light is not attracted or repelled by charged objects, so we can conclude that the beam contains small, charged particles. Because the beam was attracted to the positively charged rod and repelled from the negatively charged rod, we can further conclude that these small particles are negatively charged. His calculations showed that these newly discovered negative particles are much smaller than even the smallest atom.  Today, we call these negatively charged particles electrons.Based on his findings, Thomson proposed a model of the atom that added these electrons to Dalton's sphere model of the atom. He pictured these very small, negative electrons dispersed evenly throughout the atom.  Because scientists had observed that atoms are neutral in charge, Thomson proposed that the sphere itself is positively charged, canceling out the negative charges of the electrons. This model of the atom was accepted by the scientific community, and often called the "plum pudding model."Rutherford and the Nucleus:Ernest Rutherford designed an experiment hoping to test and support Thomson's model of the atom.  Rutherford fired heavy, fast, positively charged particles at a very thin foil made of gold. He placed a screen around the gold foil that was sensitive to the particles fired, allowing him to see how the paths of the moving particles were affected as they traveled through the thin gold foil.Based on Thomson's model of the atom, what results did Rutherford expect from this experiment? He expected the heavy, fast moving particles to be able to pass straight through the thin gold foil without being affected by the atoms of gold that made up the foil.  Since Thomson's model had the matter and mass evenly distributed throughout the atom, there would be nothing dense enough to alter the course of the heavy particles flying through the foil.What Rutherford's experiment revealed is the fact that the particles scattered instead of passing straight through was very surprising. Rutherford proposed that atoms have a small, dense nucleus that contains most of the mass of the atom.  This nucleus is what deflected the particles as they passed through the foil.  He added to this model by discovering the proton, a particle with a positive charge that is found in the nucleus.  A proton's mass is nearly 2000 times the mass of an electron, but the positive charge of one proton cancels out the negative charge of one electron.  In setting out to prove Thomson's model, Rutherford ended up proposing a major change to the atomic model.Chadwick and the Neutron:Rutherford had already discovered the nucleus and the proton, but scientists were finding that the masses of many atoms seemed to be greater than the mass of their protons. Rutherford proposed that there could be a particle with mass but no charge. He called it a neutron, and imagined it as a paired proton and electron. There was no evidence for any of these ideas, so scientists continued to study atoms. Chadwick repeated radiation experiments conducted by other scientists, such as the Curies, but with the goal of looking for the neutral particle that Rutherford described. His experiments were successful. In 1932, he was able to determine that the neutron did exist in an atom's nucleus and that its mass was very close to the mass of a proton.Current Model of the Atom:New observations and experimental results continued to change and update the model of the atom. The current model of the atom is mostly empty space.  Almost all of the atom's mass is found in the protons and neutrons, which are packed into the very small nucleus.  The electrons whiz around the nucleus in the empty space surrounding the nucleus, forming what is sometimes called the electron cloud.The drawings and models of atoms that you see in books and in this course are not to scale.  An atom this size would have a nucleus so small you could not see it in the drawing.  Even though atoms are so tiny, they are mostly empty space. If an atom had a nucleus the size of a pea, the closest electron would be two soccer fields away!

As you have seen, it can be a challenge to create a model of something that cannot be seen directly.  Scientists create models based on observations and experimental findings to help them visualize things that can be difficult to see.  In summary the atomic model changed over time:

A solid sphere John Dalton Atomic Theory Law of Multiple Proportions

A sphere with negative electrons J.J Thomson Cathode Ray Experiments Discovery of Electrons

A nucleus surrounded by electrons Ernest Rutherford Gold Foil Experiment Discovery of the Nucleus

Subatomic particles

Subatomic Particle

Location within Atom Relative Charge

Relative Mass Number

Proton The protons are found in the small nucleus at center of the atom

+1 charge 1 u The unit u stands for atomic mass unit- a very small unit of measurement used to represent the mass of atoms and subatomic particles. 1 u = 1.66 × 10-24 grams

Neutron The neutrons are found in the nucleus with the protons

No charge (neutral)

1 u A neutron has the same relative mass as a proton, and they are each about 2,000 times the mass of an electron. That means that almost all of the mass of an atom is located in the tiny nucleus.

Electron The electrons are found in the electron cloud surrounding the nucleus

-1 charge 0 u The mass of the electron is so much smaller than the masses of protons and neutrons that it is ignored when determining the relative mass of atoms.

Please get a periodic table of the elements, or visit www.ptable.com It is better if you have a printout handy.

The atomic number always tells you the number of protons in an atom, but it can also tell you the number of electrons in a neutral atom. Remember that protons have a positive (+1) charge and electrons have a negative (-1) charge. Atoms are neutral, meaning they have no net charge, in their natural state. This means that the number of electrons equals the number of protons in a neutral atom.

The atomic number of silver (Ag) is 47. This means that a neutral atom of silver has 47 protons and 47 electrons.

The atomic masses of the elements on the periodic table are measured in atomic mass units. Because the mass of an electron is so small, it is usually ignored. We can add the number of protons and neutrons together to get the mass number of an atom.

Mass Number = Number of Protons + Number of Neutrons

Use your periodic table to answer these questions Question One: A carbon atom has an atomic number of 6.

1. What are the number of protons in this atom?—62. What is this number called?—Atomic Number

Question Two: A magnesium atom has an atomic number of 12.1. Whater are the number of eletrons in this atom?—122. What is this number called?—Atomic Number

Question Three: How do you determine the Mass Number if there are thirteen neutrons in a iron (Fe) atom with an atomic number of 26?1. How to calculate the mass number?—26 + 132. What is label given to each number?—Atomic Number plus Neutrons

Question Four: Complete each of the following statements.1. Number of Protons = Atomic Number2. Number of Electrons = Atomic Number (if atom is neutral)3. Number of Neutrons = Mass Number minus Atomic Number

Question One: If a neutral hydrogen atom has an atomic number of 1 and a mass number of 1, the atom has: Protons = 1 (atomic number) Electron = 1 (atomic number) Neutrons = 0 (Neutrons = mass number − atomic number)

Question Two: If a neutral atom of chlorine has an atomic number of 17 and a mass number of 35, there are: Protons = 17 (atomic number) Electron = 17 (atomic number) Neutrons = 18 (Neutrons = mass number − atomic number)

Question Three: If a neutral fluorine atom has 9 protons and 10 neutrons, what is the atom's: Atomic Number = 9 (number of protons)

Mass Number = 19 (Mass Number = protons + neutrons) Isotopes and Weighted AveragesYou already know that the number of protons in an atom’s nucleus identifies an atom as a specific element. Atoms of the same element that have different masses with varying numbers of neutrons are called isotopes.To represent a specific isotope in writing

Question One: How many protons, neutrons, and electrons are in an atom of chlorine-37? Chlorine has an atomic number of 17. Protons = 17 (atomic number)

Electron = 17 (atomic number) Neutrons = 20 (neutrons = mass number (37)—atomic number (17)

Question Two: How many protons, neutrons, and electrons are in an atom of Br with an atomic number of 35 and a mass number of 80? Protons = 35 (atomic number) Electron = 35 (atomic number) Neutrons = 45 (neutrons = mass number (80)—atomic number (35)

Question Three: How many protons, neutrons, and electrons in an atom of beryllium-9? Beryllium has an atomic number of 4. Protons = 4 (atomic number) Electron = 4 (atomic number) Neutrons = 5 (neutrons = mass number (9)—atomic number (4)

Question Four: How many protons, neutrons, and electrons are in an atom of Ar with an atomic number of 18 and a mass number of 38? Protons = 18 (atomic number) Electron = 18 (atomic number) Neutrons = 20 (neutrons = mass number (38)—atomic number (18)

Average Atomic MassMost elements found in nature occur as a mixture of two or more isotopes. The percentages of each isotope in the mixture are the same in every sample of a given element, the average atomic mass of each element is calculated using a weighted average.The average atomic mass of an element can be calculated by multiplying the mass of each isotope by its relative abundance (the percentage represented in decimal form) and adding the results.Isotope A [percent abundance × mass] + Isotope B [percent abundance × mass] + continue for each isotope = average atomic mass

ExampleIsotope Mass

Percent Abundance(Don’t forget to change each percent value to a decimal)

Oxygen-16 16.0 u 99.762%

Oxygen-17 17.0 u 0.038%

Oxygen-18 18.0 u 0.200%

 [ .99762 × 16.0 u ] + [ .00038 × 17.0 u ] + [ .00200 × 18.0 u ] = 16.0 u

Only three significant figures, like the number you started with.

The Mole

If a sample of magnesium (Mg) atoms has a mass of 36.8 g, how many atoms are in the sample?

What would be the mass of 3.24 × 1025 atoms of oxygen (O)?

Electron Arrangement and EMR

Emission Spectrum and the Bohr Model

All systems in nature tend to take a position or arrangement that has the lowest possible potential energy. The lowest potential energy arrangement for an atom’s electrons, called the ground-state electron configuration, is the arrangement that places them closest to the nucleus.

Mapping out the ground-state arrangement of an atom’s electrons helps chemists understand many atomic properties. Electrons are involved in the formation of ions and compounds. So being able to describe the arrangement of electrons in various atoms is important in your study of chemistry.

For instance, the element hydrogen (H) is in the spot that represents the first electron of the 1s orbital. It has one electron so its electron configuration is 1s1.

Helium is in the second 1s location, representing the second electron able to be in the one orbital in the 1s sub-level. It’s electron configuration is 1s2.

The periodic table helps reveal an element’s electron configuration when you read it left to right, row by row, like you are reading this paragraph. In the interactive below, the electron configuration of carbon is determined. See how the periodic table is used to determine the electron configuration of Carbon.

Quantum Numbers