# Topic 2.

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Topic 2.

Electrons in Motion Niels Bohr ( ), a Danish scientist who worked with Rutherford, proposed that electrons must have enough energy to keep them.

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Topic 2

Atomic Structure: Basic Concepts
Topic 2 Electrons in Motion Niels Bohr ( ), a Danish scientist who worked with Rutherford, proposed that electrons must have enough energy to keep them in constant motion around the nucleus. Electrons have energy of motion that enables them to overcome the attraction of the positive nucleus.

This energy keeps the electrons moving around the nucleus.
Atomic Structure: Basic Concepts Topic 2 Electrons in Motion This energy keeps the electrons moving around the nucleus. Bohr’s view of the atom, which he proposed in 1913, was called the planetary model.

The Electromagnetic Spectrum
Atomic Structure: Basic Concepts Topic 2 The Electromagnetic Spectrum To boost a satellite into a higher orbit requires energy from a rocket motor. One way to increase the energy of an electron is to supply energy in the form of high-voltage electricity. Another way is to supply electromagnetic radiation, also called radiant energy.

The Electromagnetic Spectrum
Atomic Structure: Basic Concepts Topic 2 The Electromagnetic Spectrum Radiant energy travels in the form of waves that have both electrical and magnetic properties. These electromagnetic waves can travel through empty space, as you know from the fact that radiant energy from the sun travels to Earth every day.

The Electromagnetic Spectrum
Atomic Structure: Basic Concepts Topic 2 The Electromagnetic Spectrum As you may already have guessed, electromagnetic waves travel through space at the speed of light, which is approximately 300 million meters per second. Speed of light is represented by the lower case letter c c= (wavelength)(frequency)

The Electromagnetic Spectrum
Atomic Structure: Basic Concepts Topic 2 The Electromagnetic Spectrum Electromagnetic radiation includes radio waves that carry broadcasts to your radio and TV, microwave radiation used to heat food in a microwave oven, radiant heat used to toast bread, and the most familiar form, visible light. All of these forms of radiant energy are parts of a whole range of electromagnetic radiation called the electromagnetic spectrum.

The Electromagnetic Spectrum
Atomic Structure: Basic Concepts Topic 2 The Electromagnetic Spectrum

Wavelength and Frequency
Atomic Structure: Basic Concepts Topic 2 Wavelength and Frequency Wavelength is the measurement of a wave. A wave is measured from crest to crest or trough to trough. Frequency is the number of waves per second. 1Hz = 1 hertz =1 wave/sec = s-1 Hertz (Hz) is the SI Unit for frequency

Atomic Structure: Basic Concepts
Topic 2 Electrons and Light The spectrum of light released from excited atoms of an element is called the emission spectrum of that element.

Evidence for Energy Levels
Atomic Structure: Basic Concepts Topic 2 Evidence for Energy Levels Bohr theorized that electrons absorbed energy and moved to higher energy states. Then, these excited electrons gave off that energy as light waves when they fell back to a lower energy state.

Evidence for Energy Levels
Atomic Structure: Basic Concepts Topic 2 Evidence for Energy Levels Because electrons can have only certain amounts of energy, Bohr reasoned, they can move around the nucleus only at distances that correspond to those amounts of energy. These regions of space in which electrons can move about the nucleus of an atom are called energy levels.

The Electron Cloud Model
Atomic Structure: Basic Concepts Topic 2 The Electron Cloud Model As a result of continuing research throughout the 20th century, scientists today realize that energy levels are not neat, planetlike orbits around the nucleus of an atom. Instead, they are spherical regions of space around the nucleus in which electrons are most likely to be found.

The Electron Cloud Model
Atomic Structure: Basic Concepts Topic 2 The Electron Cloud Model Electrons themselves take up little space but travel rapidly through the space surrounding the nucleus. These spherical regions where electrons travel may be depicted as clouds around the nucleus. The space around the nucleus of an atom where the atom’s electrons are found is called the electron cloud.

The Electron Cloud Model
Atomic Structure: Basic Concepts Topic 2 The Electron Cloud Model

Electrons in Energy Level
Atomic Structure: Basic Concepts Topic 2 Electrons in Energy Level How are electrons arranged in energy levels? Each energy level can hold a limited number of electrons. The lowest energy level is the smallest and the closest to the nucleus.

Electrons in Energy Level
Atomic Structure: Basic Concepts Topic 2 Electrons in Energy Level This first energy level holds a maximum of two electrons. The second energy level is larger because it is farther away from the nucleus. It holds a maximum of eight electrons. The third energy level is larger still and holds a maximum of 18 electrons.

A hydrogen atom has only one electron. It’s in the first energy level.
Atomic Structure: Basic Concepts Topic 2 Energy Levels A hydrogen atom has only one electron. It’s in the first energy level.

Electrons in Energy Level
Atomic Structure: Basic Concepts Topic 2 Electrons in Energy Level The electrons in the outermost energy level are called valence electrons. You can also use the periodic table as a tool to predict the number of valence electrons in any atom in Groups 1, 2, 13, 14, 15, 16, 17, and 18. All atoms in Group 1, like hydrogen, have one valence electron. Likewise, atoms in Group 2 have two valence electrons.

Electrons in Energy Level
Atomic Structure: Basic Concepts Topic 2 Electrons in Energy Level An oxygen atom has eight electrons. Two of these fill the first energy level, and the remaining six are in the second energy level.

Atomic Structure: Basic Concepts
Topic 2 Lewis Dot Diagrams Because valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols.

Atomic Structure: Basic Concepts
Topic 2 Lewis Dot Diagrams A Lewis dot diagram illustrates valence electrons as dots (or other small symbols) around the chemical symbol of an element.

Each dot represents one valence electron.
Atomic Structure: Basic Concepts Topic 2 Lewis Dot Diagrams Each dot represents one valence electron. In the dot diagram, the element’s symbol represents the core of the atom—the nucleus plus all the inner electrons.

Basic Concept Questions
Topic 2 Question 1 How does the atomic number of an element differ from the element’s mass number? Answer The atomic number of an element is the number of protons in the nucleus. The mass number is the sum of the number of protons and neutrons.

Question 2 Write a Lewis dot diagram for each of the following.
Basic Concept Questions Topic 2 Question 2 Write a Lewis dot diagram for each of the following. A. Chlorine B. Calcium C. Potassium

Answer A. Chlorine B. Calcium C. Potassium Topic 2
Basic Concept Questions Topic 2 Answer A. Chlorine B. Calcium C. Potassium

Basic Concept Questions
Topic 2 Question 3 Give an example for each type of electromagnetic energy listed below. A. Ultraviolet light B. Infrared light C. Visible light

Basic Concept Questions Topic 2 Answer Sample answers: A. ultraviolet light: part of sunlight B. infrared light: radiant heat C. visible light: the spectrum of light we see as color

Energy Levels and Sublevels
Atomic Structure: Additional Concepts Topic 2 Energy Levels and Sublevels The emission spectrum for each element has a characteristic set of spectral lines. This means that the energy levels within the atom must also be characteristic of each element. But when scientists investigated multi-electron atoms, they found that their spectra were far more complex than would be anticipated by the simple set of energy levels predicted for hydrogen.

Energy Levels and Sublevels
Atomic Structure: Additional Concepts Topic 2 Energy Levels and Sublevels Notice that these spectra have many more lines than the spectrum of hydrogen.

Energy Levels and Sublevels
Atomic Structure: Additional Concepts Topic 2 Energy Levels and Sublevels Some lines are grouped close together, and there are big gaps between these groups of lines.

Energy Levels and Sublevels
Atomic Structure: Additional Concepts Topic 2 Energy Levels and Sublevels The big gaps correspond to the energy released when an electron jumps from one energy level to another.

Energy Levels and Sublevels
Atomic Structure: Additional Concepts Topic 2 Energy Levels and Sublevels The interpretation of the closely spaced lines is that they represent the movement of electrons from levels that are not very different in energy. This suggests that sublevels—divisions within a level—exist within a given energy level.

Energy Levels and Sublevels
Atomic Structure: Additional Concepts Topic 2 Energy Levels and Sublevels If electrons are distributed over one or more sublevels within an energy level, then these electrons would have only slightly different energies. The energy sublevels are designated as s, p, d, or f.

Energy Levels and Sublevels
Atomic Structure: Additional Concepts Topic 2 Energy Levels and Sublevels Each energy level has a specific number of sublevels, which is the same as the number of the energy level. For example, the first energy level has one sublevel. It’s called the 1s sublevel. The second energy level has two sublevels, the 2s and 2p sublevels

Energy Levels and Sublevels
Atomic Structure: Additional Concepts Topic 2 Energy Levels and Sublevels The third energy level has three sublevels: the 3s, 3p, and 3d sublevels; and the fourth energy level has four sublevels: the 4s, 4p, 4d, and 4f sublevels. Within a given energy level, the energies of the sublevels, from lowest to highest, are s, p, d, and f.

The Distribution of Electrons in Energy Levels
Atomic Structure: Additional Concepts Topic 2 The Distribution of Electrons in Energy Levels A specific number of electrons can go into each sublevel.

The Distribution of Electrons in Energy Levels
Atomic Structure: Additional Concepts Topic 2 The Distribution of Electrons in Energy Levels An s sublevel can have a maximum of two electrons, a p sublevel can have six electrons,

The Distribution of Electrons in Energy Levels
Atomic Structure: Additional Concepts Topic 2 The Distribution of Electrons in Energy Levels a d sublevel can have ten electrons, and an f sublevel can have 14 electrons.

Topic 2 Orbitals In the 1920s, Werner Heisenberg reached the conclusion that it’s impossible to measure accurately both the position and energy of an electron at the same time. This principle is known as the Heisenberg uncertainty principle. In 1932, Heisenberg was awarded the Nobel Prize in Physics for this discovery, which led to the development of the electron cloud model to describe electrons in atoms.

Topic 2 Orbitals The electron cloud model is based on the probability of finding an electron in a certain region of space at any given instant. In any atom, electrons are distributed into sublevels and orbitals in the way that creates the most stable arrangement; that is, the one with lowest energy.

Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Electron Configurations This most stable arrangement of electrons in sublevels and orbitals is called an electron configuration. Electrons fill orbitals and sublevels in an orderly fashion beginning with the innermost sublevels and continuing to the outermost.

Orbitals and the Periodic Table
Atomic Structure: Additional Concepts Topic 2 Orbitals and the Periodic Table The shape of the modern periodic table is a direct result of the order in which electrons fill energy sublevels and orbitals. The periodic table is divided into blocks that show the sublevels and orbitals occupied by the electrons of the atoms.

Orbitals and the Periodic Table
Atomic Structure: Additional Concepts Topic 2 Orbitals and the Periodic Table Notice that Groups 1 and 2 (the active metals) have valence electrons in s orbitals, and Groups 13 to 18 (metals, metalloids, and nonmetals) have valence electrons in both s and p orbitals.

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations Chemical properties repeat when elements are arranged by atomic number because electron configurations repeat in a certain pattern. As you move through the table, you’ll notice how an element’s position is related to its electron configuration.

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations Hydrogen has a single electron in the first energy level. Its electron configuration is 1s1. This is standard notation for electron configurations. The number 1 refers to the energy level, the letter s refers to the sublevel, and the superscript refers to the number of electrons in the sublevel.

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations Helium has two electrons in the 1s orbital. Its electron configuration is 1s2. Helium has a completely filled first energy level. When the first energy level is filled, additional electrons must go into the second energy level.

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations Lithium begins the second period. Its first two electrons fill the first energy level, so the third electron occupies the second level. Lithium’s electron configuration is 1s22s1. Beryllium has two electrons in the 2s orbital, so its electron configuration is 1s22s2 .

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations As you continue to move across the second period, electrons begin to enter the p orbitals. Each successive element has one more electron in the 2p orbitals. Carbon, for example, has four electrons in the second energy level. Two of these are in the 2s orbital and two are in the 2p orbitals. The electron configuration for carbon is 1s22s22p2.

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations At element number 10, neon, the p sublevel is filled with six electrons. The electron configuration for neon is 1s22s22p6. Neon has eight valence electrons; two are in an s orbital and six are in p orbitals.

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations Notice that neon’s configuration has an inner core of electrons that is identical to the electron configuration in helium (1s2). This insight simplifies the way electron configurations are written. Neon’s electron configuration can be abbreviated: [He]2s22p6.

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations Notice that elements in the same group have similar configurations. This is important because it shows that the periodic trends in properties, observed in the periodic table, are really the result of repeating patterns of electron configuration.

Building Electron Configurations
Atomic Structure: Additional Concepts Topic 2 Building Electron Configurations

Topic 2 Noble Gases Each period ends with a noble gas, so all the noble gases have filled energy levels and, therefore, stable electron configurations.

Topic 2 Noble Gases These stable electron configurations explain the lack of reactivity of the noble gases. Noble gases don’t need to form chemical bonds to acquire stability.

Topic 2 Transition Elements Notice in the periodic table that calcium is followed by a group of ten elements beginning with scandium and ending with zinc. These are transition elements. Now the 3d sublevel begins to fill, producing atoms with the lowest possible energy.

Topic 2 Transition Elements Like most metals, the transition elements lose electrons to attain a more stable configuration.

Inner Transition Elements
Atomic Structure: Additional Concepts Topic 2 Inner Transition Elements The two rows beneath the main body of the periodic table are the lanthanides (atomic numbers 58 to 71) and the actinides (atomic numbers 90 to 103). These two series are called inner transition elements because their last electron occupies inner-level 4f orbitals in the sixth period and the 5f orbitals in the seventh period.

Calculating Atomic Mass
Atomic Structure: Additional Concepts Topic 2 Calculating Atomic Mass

Calculating Atomic Mass
Atomic Structure: Additional Concepts Topic 2 Calculating Atomic Mass Copper exists as a mixture of two isotopes. The lighter isotope (Cu-63), with 29 protons and 34 neutrons, makes up 69.17% of copper atoms. The heavier isotope (Cu-65), with 29 protons and 36 neutrons, constitutes the remaining 30.83% of copper atoms.

Calculating Atomic Mass
Atomic Structure: Additional Concepts Topic 2 Calculating Atomic Mass The atomic mass of Cu-63 is amu, and the atomic mass of Cu-65 is amu. Use the data above to compute the atomic mass of copper.

Calculating Atomic Mass
Atomic Structure: Additional Concepts Topic 2 Calculating Atomic Mass First, calculate the contribution of each isotope to the average atomic mass, being sure to convert each percent to a fractional abundance.

Calculating Atomic Mass
Atomic Structure: Additional Concepts Topic 2 Calculating Atomic Mass The average atomic mass of the element is the sum of the mass contributions of each isotope.

Topic 2 Question 1 Write electron configurations and abbreviated electron configurations of the following elements. A. Boron B. Fluorine C. Phosphorus

Answer A. Boron B. Fluorine C. Phosphorus Topic 2
Additional Assessment Questions Topic 2 Answer A. Boron B. Fluorine C. Phosphorus

Topic 2 Question 2 The table on the next slide shows the five isotopes of germanium found in nature, the abundance of each isotope, and the atomic mass of each isotope.

Calculate the atomic mass of germanium.
Additional Assessment Questions Topic 2 Calculate the atomic mass of germanium.