Ch. 8 Chem Quiz

Describe the relationship between the properties of an element and the number of valence electrons that it contains.

-the chemical properties of elements are largely determined by the number of valence electrons they contain. -their properties are periodic because the number of valence electrons in periodic. -because element within a column in the periodic table have the same number of valence electrons, they also have similar chemical properties

What is effective nuclear charge? What is shielding?

-the effective nuclear charge is the average or net charge from the nucleus experienced by the electrons in the outermost levels. -shielding is the blocking of nuclear charge from the outermost electrons. -the shielding is primarily due to the inner electrons, although there is some interaction and shielding from the electron repulsions of the outer electrons with eachother.

What is electron affinity? What are the observed periodic trends in electron affinity?

-the electron affinity (EA) of an atom or ion is the energy charge associated with the gaining of an electron by the atom in the gaseous state. -the electron affinity is usually--though no always--negative because an atom or ion usually releases energy when it gains an electron. -The trends in electron affinity are not as regular as trends in other properties. -For main-group elements, electron affinity generally becomes more negative as you move to the right across a row in the periodic table. -There is no corresponding trend in electron affinity going down a column, with the exception of group IA which becomes more positive as you go down the column.

Explain the contributions of John Dobereiner and John Newlands to the organizations of elements according to their properties.

-the first attempt to organize the elements according to similarities in their properties was made by the German chemist Johann Dobereiner. -he grouped elements into traits: three elements with similar properties. -a more complex approach was attempted by the English chemist John Newlands. -He organized elements into octaves, analogous to musical notes. -When arranged this way, the properties of every eighth element were similar.

Describe the relationship between a main-group element's lettered group number and its valence electrons.

-the lettered group number of a main-group element is equal to the number of valence electron for that element

Which periodic property is particularly important to nerve signal transmission? Why?

-the relative size of the sodium and potassium ions is important to nerve signal transmission. -the pumps and channels within cell membranes are so sensitive that they can distinguish between the sizes of these two ions and selectively allow only one or the other to pass. -the movement of ions is the basis for the transmission of nerve signals in the brain and throughout the body.

Why are the sub levels within a principal level split into different energies for multielectron atoms but not for the hydrogen atom?

-the sub levels within a principal level split in multi electron atoms because of penetration of the outer electron region of the core electrons -the sub levels in hydrogen are not split because they are empty in the ground state.

What are periodic properties?

-a periodic property is predictable based on the element's position within the periodic table

Describe the relationship between a. the radius of a cation and that of the atom from which it forms b. the radius of an anion and that of the atom from which it forms

(a) In general, cations are much smaller than their corresponding parent. -This is because the outermost electrons are shielded from the nuclear charge in the atom and contribute greatly to the size of the atom. -When these electrons are removed to form the cation, the same nuclear charge is now acting only on the core electrons. (b) In general, anions are much larger than their corresponding atoms. -This is because the extra electrons are added to the outermost electrons but no additional protons are added to increase the nuclear charge. -The extra electrons increases the repulsions among the outermost electrons, resulting in an anion that is larger than the atom.

What is an electron configuration? Give an example.

-an electron configuration shows the particular orbitals that are occupied by electrons in an atom -some examples are H = 1s^1, He=1s^2

For transition elements, describe and explain the observed trends in atomic radius as you move: a. across the periodic table b. down a column in the periodic table

(a) The radii of transition elements stay roughly constant across each row instead of decreasing in size as in the main-group elements. -The difference is that across a row of transition elements, the number of electrons in the outermost principal energy level is nearly constant. -As another proton is added to the nucleus with each successive element, another electron is added as well, but the electron goes into a n(highest) - 1 orbital. -The number of outermost electrons stays constant, and they experience a roughly constant effective nuclear charge, keeping the radius approximately constant. (b) As you go down the first two rows of a column within the transition metals, the elements follow the same general trend in atomic radii and the main-group elements; that is, the radii get larger because you are adding outermost electrons into higher n levels.

Write a general equation for the reaction of an alkai metal with each substance. a. a halogen b. water

(a) The reactions of the alkali metals with halogens result in the formation of metal halides 2M (s) + X2 --> 2MX (s) (b) Alkali metals react with water to form the dissolved alkali metal ion, the hydroxide ion, and hydrogen gas. 2M (s) + 2H2O (l) --> 2M+ (aq) + 2OH- (aq) + H2(g)

List the number of valence electrons for each family, and explain the relationship between the number of valence electrons and the resulting chemistry of the elements in the family. a. alkali metals b. alkaline earth metals c. halogens d. oxygen family

(a) the alkali metals have one valence electron and are among the most reactive metals because their outer electron configuration. This is why the group 1A metals tend to form 1+ cations. (b) The alkaline earth metals (group 2A) have two valence electrons, have an outer electron configuration of ns^2, and also tend to be reactive metals. They lose their ns^2 electrons to form 2+ cations. (c) The halogens have seven valence electrons and an outer electron configuration of ns^2np^5. They are among the most reactive nonmetals. They are only one electron short fo a noble gas configuration and tend to react to gain that one electron, forming 1- anions. (d) The oxygen family has six valence electrons and has an outer electron configuration of ns^2np^4. They are two electrons short of a noble gas configuration and tend to react to gain those two electrons, forming 2- anions.

What are valence electrons? Why are they important?

-Balence electrons are important in chemical bonding. -For main-group elements, the valence electrons are in the outermost principal energy level. -For transition elements, we also count the outermost d electrons among the valence, even though they are not in the outermost principal energy level. -The chemical properties of an element depend on its valence electrons, which are important in bonding because they are held most loosely. -This is why the elements in a column of the periodic table have similar chemical properties: they have the same number of valence electrons

What is Coulomb's law? Explain how the potential energy of two charged particles depends on the distance between the charged particles and on the magnitude and sign of their charges.

-Coulomb's law states that the potential energy (E) of two charged particles depends on their charges (q1 and q2) and on their separation, (r). -E= 1/4πe0 * (q1q2/r) -the potential energy is positive for charges of the same spin and negative for charges of opposite spin. -the magnitude of the potential energy depends inversely on the separation between the charged particles

What are degenerate orbitals? According to Hund's rule, how are degenerate orbitals occupied?

-Degenerate orbitals are orbitals of the same energy. -In a multi electron atom, the orbitals in a sub level are degenerate. -Hund's rule states that when filling degenerate orbitals, electrons fill them singly first, with parallel spins. -This is a result of an atom's tendency to find the lowest energy state possible.

List all orbitals from 1s through 5s according to increasing energy for mulielectron atoms.

-In order of increasing energy the orbital are 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s. -The 4s orbital fill before the 3d and the 5s fill before the 4d. -They are lower in energy because of greater penetration of the 4s and 5s orbitals.

Explain the contributions of Meyer and Moseley to the periodic table.

-Meyer proposed an organization of the unknown elements based on some periodic properties -Moseley listed elements according to the atomic number rather than the atomic mass. -This resolved the problems in mendeleev's table where an increase in atomic mass did not correlate with similar properties.

Define atomic radius. For main-group elements, describe the observed trends in atomic radius as you move: a. across a period in the periodic table b. down a column in the periodic table

-One way to define atomic radii is to oncisder the distance between nonbonding atoms in molecules or atoms that are touching each other but are not bonded together. -An atomic radius determined this way to called the nonbonding atomic radius or the van der Waals radius. -The van der Waals radius represents the radius of an atom when it is not bonded to another atom. -Another way to define the size of an atom, called bonding an atomic radius or covalent radius, is defined differently for nonmetals and metas as follows: -nonmetals: one-half the distance between two of the atoms bonded together -metal: one-half the distance between two of the atoms next to each other in a crystal of the metal -A more general term, the atomic radius, refers to a set of average bonding radii determined from measurements on a large number of elements and compounds. -The atomic radius represents the radius of an atom when it is bonded to another atom and is always smaller than the van der Waals radius. (a) as you move to the right across a period in the periodic table, atomic radius decreases. (b) As you move down a column in the periodic table, atomic radius increases.

Why is electron spin important when writing electron configurations? Explains in terms of the Pauli exclusion principle.

-The pauli exclusion principle states the following: No two electrons in an atom can have the same four quantum numbers. -Becaue two electrons occupying the same orbital have three identical quantum numbers (n, l, ml), they must have different spin quantum numbers. -The Pauli exclusion principle implies that each orbital can have a maximum of only two electrons, with opposing spins

Write a general equation for the reaction of a halogen with each substance. a. a metal b. hydrogen c. another halogen

-all of the halogens are powerful oxidizing agents (a) the halogens react with metals to form metals halides. 2M(s) + nX2 --> 2MXn (s) (b) the halogens react with hydrogen to form hydrogen halides. H2 (g) + X2 --> 2HX (g) (c) The halogens react with each other to form inter halogen compounds. for example, Br2 (l) + F2 (g) --> 2BrF (g)

Describe how to write an electron configuration for a transition metal cation. Is the order of electron removal upon ionization simply the reverse of electron addition upon filling? Why or why not?

-an important exception to simply subtracting the number of electrons occurs for transition metal cations. -when writing the electron configuration of a transition metal cation, remove the electrons in the highest n-value orbitals first, even if this does not correspond to the reverse order of filling. -normally, even though the d orbital electrons add after the s orbital electrons, the s orbital electrons re lost first. -This is because (1) the ns and (n-1)d orbitals are extremely close in energy and, depending on the exact configuration, can vary in relative energy ordering and (2) as the (n-1)d orbitals begin to fill in the first transition series, the increasing nuclear charge stabilizes the (n-1)d orbitals relative to the ns orbitals. -This happens because the (n-1)d orbitals are not outermost orbitals and therefore are not effectively shields from the increases nuclear charge by the ns orbitals.

What is an orbital diagram? Provide an example.

-an orbital diagram is a different way to show the electron configuration of an atom. -it symbolizes the electron arrow in a box that represents the orbital -ex. mortal diagram for a hydrogen atom

Use the concepts of effective nuclear charge, shielding, and n value of the valence orbital to explain the trend in atomic radius as you move across a period in the periodic table.

-as you move to the right across a row in the periodic table, the n level stays the same. -however, the nuclear charge increases and the amount of shielding stays about the same because the number of inner electrons stays the same. -So the effective nuclear charge experienced by the electrons in the outermost principal energy level increases, resulting is a stronger attraction between the outermost electrons and the nucleus and, therefore, smaller atomic radii.

What are the expectations to the periodic trends in first ionization energy? Why do they occur?

-expectations occur, with elements Be, Mg, and Ca in group 2A having a higher first ionization energy than elements B, Al, and Ga in group 3A. -This expectation is caused by the change in going form the s block to the p block. -The result is that the electrons in the s orbital shield the electron in the p orbital from nuclear charge, making it easier to remove. -another expectation occurs, with N, P, and As in group 5A having a higher first ionization energy than O, S, and Se in group 6A. -this expectation is caused by the repulsion between electrons when they must occupy the same orbital. -group 5A has 3p electrons, whereas group 6A has 4p electrons. -In the group 5A elements, the p orbital are half-filled, which makes the confirmation particularly stable. -The fourth group 6A electron must pair with another electron, making it easier to remove

Which of the transition elements in the first transition series have anomalous electron configurations?

-in the first transition series of the d block, Cr and Cu have anomalous electron configurations. -Cr is expected to be [Ar]4s^23d^4, but is found to be [Ar]4s^13d^5, and Cu is expected to be [Ar]4s^23d^9, but is found to be [Ar]4s^13d^10

What is the general trend in first ionization energy as you move down column in the periodic table? As you move across a row?

-ionization energy generally decreases as you move down a column in the periodic table because electrons in the outermost principal level become farther from the positively charged nucleus and are therefore held less tightly. -ionization energy generally increases as you move to the right across a period in the periodic table because electrons in the outermost principal energy level generally experience a greater effective nuclear charge; therefore, the electrons are closer to the nucleus.

What is metallic character? What are the observed periodic trends in metallic character?

-metals are good conducts of heat and electricity, they can be pounded into flat sheets, they can be drawn into wires, they are often shiny, and they tend to lose electrons in chemical reactions. -As you move to the right across a period in the periodic table, metallic character decreases. -As you move down a column in the periodic table, metallic character increases.

What is penetration? In an atom, which electrons tend to do the most shielding (core electrons or valence electrons)?

-penetration occurs when an electrons penetrates the electrons cloud of the 1s orbital and experiences the charge of the nucleus more fully because it is less shielded by the intervening electrons. -As the outer electron undergoes penetration into the region occupied by the inner electrons, it experiences a greater nuclear charge and therefore, according to Coulomb's law, a lower energy

Copy this blank periodic table onto a sheet of paper and label each of the blocks within the table: s block, p block, d block, and f block.

-periodic table s, p, d , f sublevels

What is shielding? In an atom, which electrons tend to do the most shielding (core electrons or valence electrons)?

-shileding or screening occurs when one electron is blocked from the full effects of the nuclear charge so that the electron experiences only a part of the nuclear charge. -It is the inner (core) electrons that shield the outer electrons from the full nuclear charge.

How is the electron configuration of an anion different form that of the corresponding natural atom? How is the electron configuration of the cation different?

-the electron configuration of a main-group monatomic ion can be deduced from the electron configuration of the neutral atom and the charge of the ion. -for anions, we simply add the number of electrons required by the magnitude of the charge of the anion. -the electron configuration of cations is obtained by subtracting the number of electrons required by the magnitude of the charge.

What is ionization energy? What is the difference between first ionization energy and second ionization energy?

-the ionization energy of an atom or ion is the energy required to remove an electron from the atom or ion in the gaseous state. -the ionization energy is always positive because removing an electron take energy. -the energy required to remove the first electron is called the first ionization energy. -the energy required to remove the second electron is called the second ionization. -the second IE is always greater than the first IE

Who is credited with arranging the periodic table? How are the elements arranged in the modern periodic table?

-the modern periodic table is credited primarily to the Russian chemist Dmitri Mendeleeve. -Mendeleev's table is based on the periodic law, which states that when elements are arranged in order of increasing mass, their properties recur periodically. -Mendeleeve arranged the elements in a table in which mass increased from left to right and elements with similar properties fell in the same columns

Explain why the s block in the periodic table has only tow columns wile the p block has six.

-the number of columns in a block corresponds to the maximum number of electrons that can occupy the particular sub level of that block. -The s block has two columns corresponding to one s orbital holding a maximum of two electrons. -The p block has six column corresponding to the three p orbitals with two electrons each

The periodic table is result of the periodic law. What observations led to the periodic law? What theory explains the underlying reasons for the periodic law?

-the periodic law was based on the observations that the properties of elements recur and certain elements have similar properties -the theory that explains the existence of the periodic law is quantum mechanical theory

Describe the relationships between an element's row number in the periodic table and the highest principal quantum number in the element's electron configuration. How does this relationship differ from main-group elements, transition elements, and inner transition elements?

-the row number of main-group element is equal to the highest principal quantum number of that element. -However, the principal quantum number of the d orbital being filled across each row in the transition series is equal to the row number minus one. -For the inner transition elements, the principal quantum number of the f orbital being filled across each row in the row number minus two.

Why do the rows in the periodic table got progressively longer as you move down the table? For example, the first row contains 2 elements, the second and third rows each contain 8 elements, and the fourth and fifth rows each contain 18 elements. Explain.

-the rows in the periodic table grow progressively longer because you are adding sub levels as the n level increases

Examination of the first few successive ionization energies for a given element usually reveals a large jump between two ionization energies. For example, the successive ionization energies of magnesium show a large jump between IE2 and IE3. The successive ionization energies of aluminum show a large jump between IE3 and IE4. Explain why these jumps occur and how you might predict them.

-the second ionization energy of Mg involves remove the second outermost electron leading to an ion with a noble gas configuration for the core electrons. -the third ionization energy requires removing a core electron from an ion with a noble gas configuraiton. -this requires a tremendous amount of energy, making IE3 very high. -For Al, IE3 involves removing the third outermost electron for Al, leaving the ion with a noble gas configuration of the core electrons. -IE4 then requires removing a core electron from an ion with a noble gas configuration. -This requires a tremendous amount of energy and makes IE4 very high. -You can predict whether the IE energy is going to very high by looking for the ionization that requires removing a core electron.

Describe how to write the electron configuration for element based on its position in the periodic table.

-to use the periodic table ot write the electron configuration, find the noble gas that precedes the element -the element has the inner electron configuration of that noble gas. -place the symbol for the noble gas in [ ]. -obtain the outer electron configuration by tracing the element across the period and assigning electrons in the appropriate orbitals.