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Section 5-3: Electron Configuration and Periodic Properties

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Presentation on theme: "Section 5-3: Electron Configuration and Periodic Properties"— Presentation transcript:

1 Section 5-3: Electron Configuration and Periodic Properties
Coach Kelsoe Chemistry Pages

2 Section Objectives Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity. Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations. Define valence electrons, and state how many are present in atoms of each main-group element. Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of the main-group elements

3 Atomic Radii The best way to define the size of an atom of a certain element is to use its atomic radius. One way to express an atom’s radius is to measure the distance between the nuclei if two identical atoms are chemically bonded together.

4 Atomic Radii The atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together. The atomic radius increases as you go left and down on the periodic table.

5 Trends Trends in Clothes Trends in Shoes Trends in School
Trends in Cars Trends in Music Trends in People

6 Atomic Radii Period Trends
Notice on page 151 that there is a gradual decrease in atomic radii from left to right. If the masses increase left to right, why doesn’t the radii increase as well? The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus!

7 Atomic Radii Period Trends
The reason the atom becomes smaller is because the nucleus is positively charged and the orbitals are negatively charged. The more of a charge there is, the stronger the pull will be.

8 Atomic Radii Group Trends
In general, the atomic radii of the main-group elements increase down a group. There are a few exceptions, like aluminum and gallium. The reason is the presence of the d-block elements that precede gallium.

9 Sample Problems Of the elements magnesium, chlorine, sodium, and phosphorus, which has the largest atomic radius? Explain. Sodium, because they are all in Period 3, and atomic radii increase toward left (small Atom #). Of the elements calcium, beryllium, barium, and strontium, which has the largest atomic radius? Explain. Barium, because they are all found in Group 2, and atomic radii increase down a group.

10 Ionization Energy A + energy A + e
An electron can be removed from an atom if enough energy is supplied. Using A as a symbol for an atom of any element, the process can be expressed as follows: A + energy A + e The A+ represents an ion of element A with a single positive charge, called a 1+ ion.

11 Ionization Energy An ion is an atom or group of bonded atoms that has a positive or negative charge. Sodium, for example, forms an Na+ ion. Any process that results in the formation of an ion is called ionization. Atoms can only gain or lose electrons, not protons.

12 Ionization Energy To compare the ease with which atoms of different elements give up electrons, chemists compare ionization energies. Ionization energy is the energy required to remove one electron from a neutral atom of an element (abbreviated as IE). The energy required to remove one electron is designated IE1, to remove two electrons, IE2, and so forth.

13 Period Trends for Ionization Energy
In general, ionization energies of the main-group elements increase left-to-right across each period. The reason for the increase from left-to-right is the same reason for the decrease in atomic radii: the higher nuclear charge.

14 Group Trends for Ionization Energy
Among the main-group elements, ionization energies generally decrease down the groups. As a rule, the farther away from the nucleus the electrons are, the easier they are to lose, hence why the energies decrease down the groups.

15 Removing Electrons from Positive Ions
With enough energy, electrons can be removed from positive ions as well as from neutral atoms. The energies for removal of additional electrons from an atom are known as the second ionization energy, third ionization energy, and so on. It gets harder to lose more electrons because the nuclear force does not change.

16 Removing Electrons from Positive Ions
It requires a lot of energy to remove electrons from the Noble Gases because their outermost energy levels are full. This is why Noble Gases are generally unreactive.

17 Sample Problem Consider two main-group elements A and B. Element A has a first ionization energy of 419 KJ/mol. Element B has a first ionization energy of 1000 KJ/mol. For each element, decide if it is more likely to be in the s-block or p-block. Which element is more likely to form a positive ion? Element A has a low IE, so it will lose electrons easier so it is probably in the s-block (Positive ion) Element B is less likely to lose electrons, so it is probably in the p-block.

18 Electron Affinity Neutral atoms can also acquire electrons.
The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. Affinity is like a date or a kid’s toy…

19 Electron Affinity A + e A + energy A + e + energy A
Most atoms release energy when they acquire an electron. A + e A + energy Then again, some atoms must be forced to gain an electron by the addition of energy. A + e + energy A

20 Period Trends for Electron Affinity
Group 17 elements gain electrons most readily. The larger the negative value, the more readily it will accept electrons. In general, electron affinities become more negative across each period within the p-block. There’s an exception between Groups

21 Group Trends for Electron Affinities
Trends for electron affinities within groups are not as regular as trends for ionization energies. As a general rule, electrons add with greater difficulty down a group. This is because of the increase in nuclear charge and increase in atomic radii.

22 Adding Electrons to Negative Ions
It is always more difficult to add a second electron to an already negatively charged ion. Second IEs are always positive. It is so hard to add electrons past the full energy level state that it never happens.

23 Ionic Radii A positive ion is called a cation.
The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius. This occurs for two reasons: The smaller electron cloud because of fewer electrons. An unbalanced positive charge pulls the electrons closer.

24 Ionic Radii A negative ion is called an anion.
The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius. This occurs due to the opposite reaction of cation formation: Larger electron cloud Higher negative charge repelling

25 Valence Electrons Chemical compounds form because electrons are lost, gained, or shared between atoms. The electrons that interact in this manner are those in the highest energy levels. These are the electrons most subject to the influence of nearby atoms or ions.

26 Valence Electrons The electrons available to be lost, gained, or shared in the formation of chemical compounds are called valence electrons. Valence electrons are often located in incompletely filled main energy levels.

27 Valence Electrons For the main group elements, the valence electrons are the electrons in the outermost s and p sublevels. The inner electrons are in filled energy levels and are held too tightly by the nucleus to be involved in compound formation. Group 1 elements have one valence electron. Group 2 elements have two. Groups have valence electrons equal to the group number minus 10. (Group 14 = 4)

28 Electronegativity The valence electrons hold atoms together in chemical compounds. Some compounds have a concentrated negative charge closer to one atom than the other, which has an effect on the chemical properties of that compound.

29 Electronegativity Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons. Fluorine is the most electronegative element, with an electronegativity value of 4.

30 Trends for Electronegativity
Electronegativities tend to increase across each period, but there are exceptions. Alkali and alkali-earth metals are the least electronegative. Electronegativities tend to either decrease down a group or remain about the same. The combination of period and group trends result in the highest values belonging to the elements in the upper right corner of the table.

31 Electron Configurations for Cations and Anions

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