Atoms the building blocks of matter
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Atoms: The Building Blocks of Matter. Foundations of Atomic Theory. Nearly all chemists in late 1700s accepted the definition of an element as a substance that cannot be broken down further Knew about chemical reactions

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Atoms: The Building Blocks of Matter

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Atoms: The Building Blocks of Matter

Foundations of Atomic Theory

  • Nearly all chemists in late 1700s accepted the definition of an element as a substance that cannot be broken down further

  • Knew about chemical reactions

  • Great disagreement as to whether elements always combine in the same ratio when forming a specific compound

Law of Conservation of Mass

  • With the help of improved balances, investigators could accurately measure the masses of the elements and compounds they were studying

  • This lead to discovery of several basic laws

  • Law of conservation of mass states that mass is neither destroyed nor created during ordinary chemical reactions or physical changes

Law of Definite Proportions

  • law of definite proportions  A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound

Each of the salt crystals shown here contains exactly

39.34% sodium and 60.66% chlorine by mass.

Law of Multiple Proportions

  • Two elements sometimes combine to form more than one compound

    • For example, the elements carbon and oxygen form two compounds, carbon dioxide and carbon monoxide

  • Consider samples of each of these compounds, each containing 1.0 g of carbon

  • In carbon dioxide, 2.66 g of oxygen combine with 1.0 g of carbon

  • In carbon monoxide, 1.33 g of oxygen combine with 1.0 g of carbon

  • The ratio of the masses of oxygen in these two compounds is exactly 2.66 to 1.33, or 2 to 1

1808 John Dalton

  • Proposed an explanation for the law of conservation of mass, the law of definite proportions, and the law of multiple proportions

  • He reasoned that elements were composed of atoms and that only whole numbers of atoms can combine to form compounds

  • His theory can be summed up by the following statements

Dalton’s Atomic Theory

  • All matter is composed of extremely small particles called atoms.

  • Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.

  • Atoms cannot be divided, created or destroyed.

  • Atoms of different elements combine in simple whole-number ratios to form chemical compounds.

  • In chemical reactions, atoms are combined, separated, or rearranged.

Modern Atomic Theory

  • Dalton turned Democritus’sidea into a scientific theory which was testable

  • Not all parts of his theory have been proven correct

  • Ex. We know atoms are divisible into even smaller particles

  • We know an element can have atoms with different masses

Structure of the Atom

All atoms consist of two regions

  • Nucleus very small region located near the center of an atom

    • In the nucleus there is at least one positively charged particle called the proton

    • Usually at least one neutral particle called the neutron

  • Surrounding the nucleus is a region occupied by negatively charged particles called electrons

Discovery of the Electron

  • Resulted from investigations into the relationship between electricity and matter

  • Late 1800s, many experiments were performed  electric current was passed through different gases at low pressures

  • These experiments were carried out in glass tubes known as cathode-ray tubes

Cathode Rays and Electrons

  • Investigators noticed that when current was passed through a cathode-ray tube, the opposite end of the tube glowed

  • Hypothesized that the glow was caused by a stream of particles, which they called a cathode ray

  • The ray traveled from the cathode to the anode when current was passed through the tube


  • 1. cathode rays deflected by magnetic field in same way as wire carrying electric current (known to have negative charge)

  • 2. rays deflected away from negatively charged object

  • Observations led to the hypothesis that the particles that compose cathode rays are negatively charged

  • Strongly supported by a series of experiments carried out in 1897 by the English physicist Joseph John Thomson

  • He was able to measure the ratio of the charge of cathode-ray particles to their mass

  • He found that this ratio was always the same, regardless of the metal used to make the cathode or the nature of the gas inside the cathode-ray tube

  • Thomson concluded that all cathode rays are composed of identical negatively charged particles, which were later named electrons

Charge and Mass of the Electron

  • Confirmed that the electron carries a negative electric charge

  • Because cathode rays have identical properties regardless of the element used to produce them, it was concluded that electrons are present in atoms of all elements

  • Cathode-ray experiments provided evidence that atoms are divisible and that one of the atom’s basic constituents is the negatively charged electron

  • Thomson’s experiment revealed that the electron has a very large charge for its tiny mass

  • Mass of the electron is about one two-thousandth the mass of the simplest type of hydrogen atom (the smallest atom known)

  • Since then found that the electron has a mass of 9.109 × 10−31 kg, or 1/1837 the mass of the hydrogen atom

  • Based on information about electrons, two inferences made about atomic structure

  • b/c atoms are neutral they must have positive charge to balance negative electrons

  • b/c electrons have very little mass, atoms must have some other particles that make up most of the mass

Thomson’s Atom

  • Plum pudding model (based on English dessert)

  • Negative electrons spread evenly through positive charge of the rest of the atom

  • Like seeds in a watermelon

Discovery of Atomic Nucleus

  • 1911 by New Zealander Ernest Rutherford and his associates Hans Geiger and Ernest Marsden

  • Bombarded a thin, gold foil with fast-moving alpha particles (positively charged particles with about four times the mass of a hydrogen atom)

  • Assume mass and charge were uniformly distributed throughout atoms of gold foil (from Thomson’s model of the atom)

  • Expected alpha particles to pass through with only slight deflection

What Really Happened…

  • Most particles passed with only slight deflection

  • However, 1/8,000 were found to have a wide deflection

  • Rutherford explained later it was “as if you have fired a 15-inch artillery shell at a piece of tissue paper and it came back and hit you.”


  • After 2 years, Rutherford finally came up with an explanation

  • The rebounded alpha particles must have experienced some powerful force within the atom

  • The source of this force must occupy a very small amount of space because so few of the total number of alpha particles had been affected by it

  • The force must be caused by a very densely packed bundle of matter with a positive electric charge

  • Rutherford called this positive bundle of matter the nucleus

  • Rutherford had discovered that the volume of a nucleus was very small compared with the total volume of an atom

  • If the nucleus were the size of a marble, then the size of the atom would be about the size of a football field

  • But where were the electrons?

  • Rutherford suggested that the electrons surrounded the positively charged nucleus like planets around the sun

  • He could not explain, however, what kept the electrons in motion around the nucleus

Rutherford’s Atom

Composition of Atomic Nucleus

  • Except hydrogen, all atomic nuclei made of two kinds of particles

    • Protons

    • Neutrons

  • Protons = positive

  • Neutrons = neutral

  • Electrons = negative

  • Atoms are electrically neutral, so number of protons and electrons IS ALWAYS THE SAME

  • The nuclei of atoms of different elements differ in the number of protons they contain and therefore in the amount of positive charge they possess

  • So the number of protons in an atom’s nucleus determines that atom’s identity

Forces in the Nucleus

  • Usually, particles that have the same electric charge repel one another

  • Would expect a nucleus with more than one proton to be unstable

  • When two protons are extremely close to each other, there is a strong attraction between them

  • Nuclear forces short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together

The Sizes of atoms

  • Area occupied by electrons is electron cloud – cloud of negative charge

  • Radius of atom is distance from center of nucleus to outer portion of cloud

  • Unit – picometer (10-12 m)

  • Atomic radii range from 40-270 pm

  • Very high densities – 2 x 108 tons/cm3

Counting Atoms

Atomic Number

  • Atoms of different elements have different numbers of protons

  • Atomic number (Z) number of protons in the nucleus of each atom of that element

  • Shown on periodic table

  • Atomic number identifies an element







  • Hydrogen and other atoms can contain different numbers of neutrons

  • Isotope atoms of same element that have different masses

Isotopes of Hydrogen

Mass Number

  • Mass number total number of protons and neutrons in the nucleus of an isotope

Sample Problem

  • How many protons, electrons, and neutrons are there in an atom of chlorine-37?

  • atomic number = number of protons = number of electrons

  • mass number = number of neutrons + number of protons

  • mass number of chlorine-37 − atomic number of chlorine = number of neutrons in chlorine-37

  • mass number − atomic number = 37 (protons plus neutrons) − 17 protons = 20 neutrons

Practice Problems

  • How many protons, electrons, and neutrons are in an atom of bromine-80?

  • Answer 35 protons, 35 electrons, 45 neutrons

2. Write the nuclear symbol for carbon-13.

  • Answer 136C

3. Write the hyphen notation for the element that contains 15 electrons and 15 neutrons.

  • Answer phosphorus-30

Relative Atomic Masses

  • Because atoms are so small, scientists use a standard to control the units of atomic mass  carbon-12

  • Randomly assigned a mass of exactly 12 atomic mass units (amu)

  • 1 amu is exactly 1/12 the mass of a carbon-12 atom

  • Atomic mass of any atom determined by comparing it with mass of C-12

  • Ex. H  1/12 C-12, so 1 amu

Average Atomic Masses

  • Average atomic mass the weighted average of the atomic masses of the naturally occurring isotopes of an element

  • Ex. You have a box containing 2 size of marbles

    • 25% are 2.00g each

    • 75% are 3.00g each

      25% = 0.2575% = 0.75

      (2.00g x 0.25) + (3.00g x 0.75) = 2.75g

Calculating Average Atomic Mass


69.17% Cu-63 – 62.929599 amu

30.83% Cu-65 – 64.927793 amu

(0.6917 x 62.929599) + (0.3083 x 64.927793)

= 63.55 amu

Relating Mass to Numbers of Atoms

  • The relative atomic mass scale makes it possible to know how many atoms of an element are present in a sample of the element with a measurable mass

  • Three very important concepts provide the basis for relating masses in grams to numbers of atoms

  • The mole

  • Avogadro’s number

  • Molar mass

The Mole

  • SI unit for an amount of substance (like 1 dozen = 12)

  • Mole (mol)  amount of a substance that contains as many particles as there are atoms in exactly 12g of C-12

Avogadro’s Number

  • Avogadro’s number the number of particles in exactly one mole of a pure substance

  • 6.022 x 1023

  • How big is that?

  • If 5 billion people worked to count the atoms in one mole of an element, and if each person counted continuously at a rate of one atom per second, it would take about 4 million years for all the atoms to be counted

Molar Mass

  • Molar mass the mass of one mole of a pure substance

  • Written in unit g/mol

  • Found on periodic table (atomic mass)

  • Ex. Molar mass of H = 1.008 g/mol


How many grams of helium are there in 2 moles of helium?

2.00 mol He x = ? g He

2.00 mol He x 4.00 g He= 8.00 g He

1 mole He

Practice Problems

What is the mass in grams of 3.50 mol of the element copper, Cu?

222 g Cu

What is the mass in grams of 2.25 mol of the element iron, Fe?

126 g Fe

What is the mass in grams of 0.375 mol of the element potassium, K?

14.7 g K

What is the mass in grams of 0.0135 mol of the element sodium, Na?

0.310 g Na

What is the mass in grams of 16.3 mol of the element nickel, Ni?

957 g Ni

A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced?

0.441 mol Al

How many moles of calcium, Ca, are in 5.00 g of calcium?

0.125 mol Ca

How many moles of gold,Au, are in 3.60 × 10−10 g of gold?

1.83 × 10−12 mol Au

Conversions with Avogadro’s Number

How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver?

Given: 3.01 × 1023 atoms of Ag

Unknown: amount of Ag in moles

Ag atoms ×= moles Ag

3.01x1023atomsAg x =0.500 mol Ag

Practice problems

How many moles of lead, Pb, are in 1.50 × 1012 atoms of lead?

2.49 × 10−12 mol Pb

How many moles of tin, Sn, are in 2500 atoms of tin?

4.2 × 10−21 mol Sn

How many atoms of aluminum, Al, are in 2.75 mol of aluminum?

1.66 × 1024 atoms Al

What is the mass in grams of 7.5 × 1015 atoms of nickel, Ni?

7.3 × 10−7 g Ni

How many atoms of sulfur, S, are in 4.00 g of sulfur?

7.51 × 1022 atoms S

What mass of gold,Au, contains the same number of atoms as 9.0 g of aluminum,Al?

66 g Au

Arrangement of Electrons in Atoms

  • Visible light is a kind of electromagnetic radiation form of energy that exhibits wavelike behavior as it travels through space

    • X-rays, UV, infrared, etc.

  • Together, all forms of electromagnetic radiation form the electromagnetic spectrum

Properties of EM Radiation

  • All forms of EM radiation travel at a constant speed (3.0 × 108 m/s) through a vacuum and at slightly slower speeds through matter

  • It has a repetitive nature, which can be described by the measurable properties  wavelength and frequency

  • Wavelength (λ)  the distance between corresponding points on neighboring waves

  • Frequency (ν)  the number of waves that pass a given point in a specific time, usually one second

The Photoelectric Effect

  • Early 1900s, scientists conducted two experiments involving relations of light and matter that could not be explained by the wave theory of light

  • One experiment involved a phenomenon known as the photoelectric effect

  • Photoelectric effect  the emission of electrons from a metal when light shines on the metal

The Mystery

  • For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum—regardless of how long the light was shone

  • Light was known to be a form of energy, capable of knocking loose an electron from a metal

  • Wave theory of light predicted light of any frequency could supply enough energy to eject an electron

  • Couldn’t explain why light had to be a minimum frequency in order for the photoelectric effect to occur

The Particle Description of Light

  • Explanation  1900, German physicist Max Planck was studying the emission of light by hot objects

  • Proposed a hot object does not emit electromagnetic energy continuously, as would be expected if the energy emitted were in the form of waves

  • Instead, Planck suggested the object emits energy in small, specific amounts called quanta

  • Quantum  the minimum quantity of energy that can be lost or gained by an atom

  • Planck projected the following relationship between a quantum of energy and the frequency of radiation

    E = hν

  • In the equation,

    • E is the energy, in joules, of a quantum of radiation

    • ν is the frequency of the radiation emitted

    • h is a constant now known as Planck’s constant

      • h = 6.626 × 10−34 J• s

1905 Albert Einstein

  • Expanded on Planck’s theory by introducing the idea that electromagnetic radiation has a dual wave-particle nature

  • Light displays many wavelike properties, it can also be thought of as a stream of particles

  • Einstein called these particles photons

  • Photon aparticle of electromagnetic radiation having zero mass and carrying a quantum of energy

  • The energy of a specific photon depends on the frequency of the radiation

    Ephoton = hν

The Hydrogen-Atom Line-Emission Spectrum

  • When current is passed through a gas at low pressure, the potential energy of some of the gas atoms increases

  • The lowest energy state of an atom is its ground state

  • A state in which an atom has a higher potential energy than it has in its ground state is an excited state

Returning to Ground State

  • When an excited atom returns to its ground state, it gives off the energy it gained in the form of electromagnetic radiation

  • Example of this process is a neon sign

Excited neon atoms emit light when falling back to the ground state or to a lower-energy excited state.

  • Whenever an excited hydrogen atom falls back from an excited state to its ground state or to a lower-energy excited state, it emits a photon of radiation

  • The energy of this photon (Ephoton = hν) is equal to the difference in energy between the atom’s initial state and its final state

  • Changes of energy (transition of an electron from one orbit to another) is done in isolated quanta

  • Quanta are not divisible

  • There is sudden movement from one specific energy level to another, with no smooth transition

  • There is no ``in-between‘‘

  • Hydrogen atoms emit only specific frequencies of light  showed that energy differences between the atoms’ energy states were fixed

  • The electron of a hydrogen atom exists only in very specific energy states

Bohr Model of the Hydrogen Atom

  • Solved in 1913 by the Danish physicist Niels Bohr

  • Proposed a model of the hydrogen atom that linked the atom’s electron with photon emission

  • According to the model, the electron can circle the nucleus only in allowed paths, or orbits

  • When the electron is in one of these orbits, the atom has a definite, fixed energy

  • The electron, and therefore the hydrogen atom, is in its lowest energy state when it is in the orbit closest to the nucleus

  • This orbit is separated from the nucleus by a large empty space where the electron cannot exist

  • The energy of the electron is higher when it is in orbits farther from the nucleus

The Rungs of a Ladder

  • The electron orbits or atomic energy levels in Bohr’s model can be compared to the rungs of a ladder

  • When you are standing on a ladder, your feet are on one rung or another

  • The amount of potential energy that you possess relates to standing on the first rung, the second rung…

  • Your energy cannot relate to standing between two rungs because you cannot stand in midair

  • In the same way, an electron can be in one orbit or another, but not in between

Moving up in the Orbit…

  • Explaining spectral lines:

  • While in a given orbit, electron isn’t gaining or losing energy

  • Can move to higher-energy orbit by gaining energy equal to difference in energy between higher-energy orbit and initial lower-energy orbit

  • When hydrogen is excited, electron is in higher-energy orbit

  • When electron falls to lower energy level, a photon is emitted in process called emission

  • Photon’s energy is equal to energy difference between initial high energy and final lower energy level

  • To move electron from low energy to high energy level, energy must be added to atom in process called absorption

  • When a hydrogen atom is in an excited state, its electron is in a higher-energy orbit

  • When the atom falls back from the excited state, the electron drops down to a lower-energy orbit

  • In the process, a photon is emitted that has an energy equal to the energy difference between the initial higher-energy orbit and the final lower-energy orbit

Electrons as Waves

  • Photoelectric effect and H spectrum showed light behaving as wave and particle

  • Could electrons have dual wave-particle nature?

  • 1924 Louis de Broglie proposed answer that revolutionized our understanding of matter

  • DeBroglie pointed out behavior of electrons in Bohr’s orbits similar to known behavior of waves

  • Any wave in limited space has only certain frequencies

  • Electrons should be considered as being in limited space around the nucleus

  • So electron waves could exist only at specific frequencies

  • According to E = hv, frequencies will correspond to specific energies – the quantized energies of Bohr’s orbits

The Heisenberg Uncertainty Principle

  • If electrons are both particles and waves, then where are they in the atom?

  • To answer this question, it is important to consider a proposal first made in 1927 by the German theoretical physicist Werner Heisenberg

  • Heisenberg’s idea involved detection of electrons by interaction with photons

  • Photons have about same energy as electrons

  • Trying to locate specific electron with photon knocks electron off course

  • So there is always uncertainty trying to locate an electron

  • Heisenberg uncertainty principlestates that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle

Schrodinger Wave Equation

  • Used hypothesis that electrons have dual nature to develop equation that treated electrons in atoms as waves

  • Bohr assumed quantization was fact

  • Schrodinger didn’t assume – it was a natural result of his equation

  • Only waves of specific energies (frequencies) provided solutions to equation

  • With Heisenberg uncertainty, wave equation laid foundation for modern quantum theory

  • Quantum theory – describes mathematically the wave properties of electrons and other small particles


  • Electrons do not travel around the nucleus in neat orbits

  • Instead, they exist in certain regions called orbitals

  • Orbital  a three-dimensional region around the nucleus that indicates the probable location of an electron

Atomic Orbitals and Quantum Numbers

  • Bohr’s model – electrons of increasing energy occupy orbits farther and farther from nucleus

  • Schrodinger equation – electrons in orbitals have quantized energies

  • Energy level is not the only characteristic of an orbital shown by solving Schrodinger equation

Atomic Orbitals and Quantum Numbers

  • In order to completely describe orbitals, scientists use quantum numbers

  • Quantum numbers  specify the properties of atomic orbitals and theproperties of electrons in orbitals

  • First 3 numbers result from solutions to Schrodinger’s equation

  • Indicate energy level, shape, and orientation of each orbital

The Quantum Numbers

  • Principal quantum number  the main energy level occupied by the electron

  • Angular momentum quantum number  the shape of the orbital

  • Magnetic quantum number  the orientation of the orbital around the nucleus

  • Spin quantum number  the two fundamental spin states of an electron in an orbital

Principal Quantum Number

  • Symbolized by n

  • Indicates the main energy level occupied by the electron

  • Values are positive integers only

  • As n increases, the electron’s energy and its average distance from the nucleus increase

Angular Momentum Quantum Number

  • Except at the first main energy level, orbitals of different shapes— known as sublevels—exist for a given value of n

  • The angular momentum quantum number, symbolized by l, indicates the shape of the orbital

  • The values of l allowed are zero and all positive integers less than or equal to n − 1

    • n = 2 can have one of two shapes

    • l = 0 and l = 1

n = 1

  • l = 0 only

  • For l = 0, the corresponding letter designation for the orbital shape is s

  • Shape of s is a sphere

n = 2

  • l = 0 (s), l = 1 (p)

  • For l = 1, , the corresponding letter designation for the orbital shape is p

  • Shape of p is

n = 3

  • l = 0 (s), l = 1 (p), l = 2 (d)

  • For l = 2, the corresponding letter designation for the orbital shape is d

  • Shape of d is

n= 4

  • l = 0 (s),1 (p), 2 (d), 3 (f)

  • For l = 3, the corresponding letter designation for the orbital shape is f

  • Shape is too complicated to picture

  • Each atomic orbital is designated by the principal quantum number followed by the letter of the sublevel

  • For example, the 1s sublevel is the s orbital in the first main energy level, while the 2p sublevel is the set of p orbitals in the second main energy level

Magnetic Quantum Number

  • Atomic orbitals can have the same shape but different orientations around the nucleus

  • The magnetic quantum number, symbolized by m, indicates the orientation of an orbital around the nucleus

S orbital

  • Because an s orbital is spherical and is centered around the nucleus, it has only one possible orientation

  • This orientation corresponds to a magnetic quantum number of m = 0

  • There is therefore only one s orbital in each s sublevel

P orbital

  • p orbital can extend along the x, y, or z axis

  • There are therefore three porbitals in each p sublevel, which are designated as px, py, and pzorbitals

  • The three porbitals occupy different regions of space and correspond, in no particular order, to values of m = −1, m = 0, and m = +1

D orbital

  • There are five different d orbitals in each d sublevel

  • The five different orientations, including one with a different shape, correspond to values of m = −2, m = −1, m = 0, m = +1, and m = +2

Spin Quantum Number

  • An electron in an orbital can be thought of as spinning on an axis

  • It spins in one of two possible directions, or states

  • The spin quantum number has only two possible values — +1/2 , −1/2

  • Indicate the two fundamental spin states of an electron in an orbital

  • ***A single orbital can hold a maximum of two electrons, which must have opposite spins***

Electron Configurations

  • Quantum model of atom better than Bohr model b/c it describes the arrangements of electrons in atoms other than hydrogen

  • Electron configuration the arrangement of electrons in an atom

  • b/c atoms of different elements have different numbers of e-, the e- configuration for each element is unique

  • Electrons in atoms assume arrangements that have lowest possible energies

  • Ground-state electron configuration lowest-energy arrangement of electrons for each element

  • Some rules combined with quantum numbers let us determine the ground-state e- configurations

Rules Governing Electron Configurations

  • To build up electron configurations for the ground state of any particular atom, first the energy levels of the orbitals are determined

  • Then electrons are added to the orbitals one by one according to three basic rules

  • Aufbau principle

  • Pauli exclusion principle

  • Hund’s rule

Aufbau Principle

  • The first rule shows the order in which electrons occupy orbitals

  • According to the Aufbau principle, an electron occupies the lowest-energy orbital that can receive it

  • The orbital with the lowest energy is the 1s orbital

  • Ground-state hydrogen atom  electron in this orbital

  • The 2s orbital is the next highest in energy, then the 2p orbitals

  • Beginning with the third main energy level, n = 3, the energies of the sublevels in different main energy levels begin to overlap

  • The 4s sublevel is lower in energy than the 3d sublevel

  • Therefore, the 4s orbital is filled before any electrons enter the 3d orbitals

  • (Less energy is required for two electrons to pair up in the 4s orbital than for a single electron to occupy a 3d orbital)

Pauli Exclusion Principle

  • The second rule reflects the importance of the spin quantum number

  • According to the Pauli exclusion principle, no two electrons in the same atom can have the same set of four quantum numbers

Hund’s Rule

  • The third rule requires placing as many unpaired electrons as possible in separate orbitals in the same sublevel

  • Electron-electron repulsion is minimized  the electron arrangements have the lowest energy possible

  • Hund’s rule orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin

Orbital Notation

Electron-Configuration Notation

  • Electron-configuration notation eliminates the lines and arrows of orbital notation

  • Instead, the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation

  • The hydrogen configuration is represented by 1s1

  • The superscript indicates that one electron is present in hydrogen’s 1s orbital

  • The helium configuration is represented by 1s2

  • Here the superscript indicates that there are two electrons in helium’s 1s orbital

Sample Problem

The electron configuration of boron is 1s22s22p1. How many electrons are present in an atom of boron? What is the atomic number for boron? Write the orbital notation for boron.

Practice Problems

  • The electron configuration of nitrogen is 1s22s22p3. How many electrons are present in a nitrogen atom? What is the atomic number of nitrogen? Write the orbital notation for nitrogen.

Practice Problems

The electron configuration of fluorine is 1s22s22p5. What is the atomic number of fluorine? How many of its p orbitals are filled? How many unpaired electrons does a fluorine atom contain?

  • Answer

  • 9

  • 2

  • 1

Elements of the Second Period

  • According to Aufbau principle, after 1s orbital is filled, the next electron occupies the s sublevel in the second main energy level

  • Lithium, Li, has a configuration of 1s22s1

  • The electron occupying the 2s level of a lithium atom is in the atom’s highest, or outermost, occupied level

  • The highest occupied level  the electron-containing main energy level with the highest principal quantum number

  • The two electrons in the 1s sublevel of lithium are no longer in the outermost main energy level

  • They have become inner-shell electronselectrons that are not in the highest occupied energy level

  • The fourth electron in an atom of beryllium, Be, must complete the pair in the 2s sublevel because this sublevel is of lower energy than the 2p sublevel

  • With the 2s sublevel filled, the 2p sublevel, which has three vacant orbitals of equal energy, can be occupied

  • One of the three p orbitals is occupied by a single electron in an atom of boron, B

  • Two of the three p orbitals are occupied by unpaired electrons in an atom of carbon, C

  • And all three p orbitals are occupied by unpaired electrons in an atom of nitrogen, N

Elements of Third Period

  • After the outer octet is filled in neon, the next electron enters the s sublevel in the n = 3 main energy level

  • Atoms of sodium, Na, have the configuration 1s22s22p63s1

  • Once past Neon, can use Noble-gas notation

Noble-Gas Notation

  • The Group 18 elements (helium, neon, argon, krypton, xenon, and radon) are called the noble gases

  • To simplify sodium’s notation, the symbol for neon, enclosed in square brackets, is used to represent the complete neon configuration:

    [Ne] = 1s22s22p6

  • This allows us to write sodium’s electron configuration as [Ne]3s1, which is called sodium’s noble-gas notation

Elements of the Fourth Period

  • Period begins by filling 4s orbital (empty orbital of lowest energy)

  • First element in fourth period is potassium, K

  • E- configuration [Ar]4s1

  • Next element is calcium, Ca

  • E- configuration [Ar]4s2

  • With 4s sublevel filled, 4p and 3d sublevels are next available empty orbitals

  • Diagram shows 3d sublevel is lower energy than 4p sublevel

  • So 3d is filled next

  • Total of 10 electrons can fill 3d orbitals

  • These are filled from element scandium (Sc) to zinc (Zn)

  • Scandium has e- configuration [Ar]3d14s2

  • Titanium, Ti, has configuration [Ar]3d24s2

  • Vanadium, V, has configuration [Ar]3d34s2

  • Up to this point, 3 e- with same spin added to 3 separate d orbitals (required by Hund’s rule)

  • Chromium, Cr, has configuration [Ar]3d54s1

  • We added one electron to 4th 3d orbital

  • We also took a 4s electron and added it to 5th 3d orbital

  • Seems to be against Aufbau principle

  • In reality, [Ar]3d54s1 is lower energy than [Ar]3d44s2

  • Having 6 outer orbitals with unpaired electrons in 3d orbital is more stable than having 4 unpaired electrons in 3d orbitals and forcing 2 electrons to pair in 4s

  • For tungsten, W, (same group as chromium) having 4 e- in 5d orbitals and 2 e- paired in 6s is most stable arrangement

  • No easy explanantion

  • Manganese, Mn, has configuration [Ar]3d54s2

  • Added e- goes to 4s orbital, completely filling it and leaving 3d half-filled

  • Starting with next element, e- continue to pair in d orbitals

  • So iron, Fe, has configuration [Ar]3d64s2

  • Cobalt, Co, has [Ar]3d74s2

  • Nickel, Ni, has [Ar]3d84s2

  • With copper, Cu, an electron from 4s moves to pair with e- in 5th 3d orbital

  • Configuration: [Ar]3d104s1 (most stable)

  • For zinc, Zn, 4s sublevel filled to give [Ar]3d104s2

  • For next elements, 1 e- added to 4p orbitals according to Hund’s rule

Elements of Fifth Period

  • In 18 elements of 5th period, sublevels fill in similar way to 4th period elements

  • BUT they start at 5s level instead of 4s level

  • Fill 5s, then 4d, and finally 5p

  • There are exceptions just like in 4th period

Practice Problem

Write both the complete electron-configuration notation and the noble-gas notation for iron, Fe.

Write both the complete electron configuration notation and the noble-gas notation for iodine, I. How many inner-shell electrons does an iodine atom contain?

  • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p65s2 4d10 5p5

  • [Kr]4d10 5s2 5p5

  • 46

How many electron-containing orbitals are in an atom of iodine? How many of these orbitals are filled? How many unpaired electrons are there in an atom of iodine?

  • 27

  • 26

  • 1

Write the noble-gas notation for tin, Sn. How many unpaired electrons are there in an atom of tin?

  • [Kr] 5s2 4d10 5p2

  • 2

Without consulting the periodic table or a table in this chapter, write the complete electron configuration for the element with atomic number 25.

  • 1s2 2s2 2p6 3s2 3p6 4s2 3d5

Elements of 6th and 7th periods

  • 6th period has 32 elements

  • To build up e- configurations for elements in this period, e- first added first to 6s orbital in cesium, Cs, and barium, Ba

  • Then in lanthanum, La, e- added to 5d orbital

  • With next element, cerium, Ce, 4f orbitals begin to fill

  • Ce: [Xe]4f15d16s2

  • In next 13 elements, 4f orbitals filled

  • Next 5d orbitals filled

  • Period finished by filling 6p orbitals

  • (some exceptions happen)

Practice Problem

Write both the complete electron-configuration notation and the noble-gas notation for a rubidium atom.

  • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1

  • [Kr]5s1

Write both the complete electron configuration notation and the noble-gas notation for a barium atom.

  • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2

  • [Xe]6s2

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