chapter 2: periodic properties: atoms and simple...

UBC CHEM 121 Ethan Lee

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Chapter 2: Periodic Properties: Atoms and Simple Compounds

Atomic Structure -Atoms mainly consist of electrons, protons, and neutrons.

Particle Charge Mass (u)

proton +1 1.0073

neutron 0 1.0087

electron -1 0.0005486

1 atomic mass unit (u) →

the mass of one atom of carbon-12

-The nucleus is composed of protons and neutrons whereas the electrons are located outside of the nucleus. -The nucleus is tiny compared to the rest of the atom. -Atoms are mostly empty space. -It is the electrons of an atom that are involved in chemical reactions.

-The identity of a chemical element is determined by the # of protons in the nucleus. ↘called the atomic # (Z) -protons + neutrons = mass #

-You usually don’t need to write atomic #.


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isotope → a different form of an element with the same amount of protons but a different amount of neutrons -Isotopes of H:

1H 2H 3H

1 proton 1 proton 1 proton

0 neutrons 1 neutron 2 neutrons

ex. 1H2O → 18g/mol 2H2O → 20g/mol -In nature, the isotopes of H are not found in equal amounts. -Many elements are composed of a certain ratio of isotopes. natural abundance → the % of each naturally occurring sample of an element ex. The natural abundance of 1H is 99.9%, 2H is 0.0156%, and 3H is about 10-18 atoms per 1H atom. -Isotopes of the same element will react in similar ways due to having the same # of electrons, which are the particles involved in chemical reactions. -The mass # on the periodic tale can tell you which isotope has a higher natural abundance. ex. boron’s atomic weight = 10.811 -The two naturally occurring isotopes are 10B and 11B. -10.811 is closer to 11 so the % of 11B is higher. Atomic Weight atomic weight → the weighted average of the masses of the naturally occurring isotopes

ex. The atomic weight of carbon is 12.011u (close to 12u) because in a naturally occurring sample of carbon, 98.93% of the atoms are 12C and 1.07% are 13C.

0.9893 x 12 = 11.872 0.0107 x 13.003 = 0.13913

12.011

ex. A naturally occurring sample of copper consists of two isotopes, 63Cu = 62.9295u and 65Cu = 64.9278u.

-Using the atomic weight (AW) of copper from the periodic table, the natural abundance of each isotope can be found:

AW = 63.546u

AW = (% 63Cu)(mass 63Cu) + (% 65Cu)(mass 65Cu)

Let a = % 63Cu and b = % 65Cu


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63.546u = a(62.9295u) + b(64.9278u) % 63Cu + % 65Cu = 100%, so a + b = 1 and b = 1 - a

63.546 = 62.929a + 64.9278(1 - a)

a = 69.15%

Therefore, the % 63Cu is 69.15% and that of 65Cu is 30.85%.

Periodic Table of the Elements -The 117 known elements are arranged into 18 groups according to similarities in behavior. -Hydrogen is unique in that it is not easily classified into a specific group as it has chemical properties similar to both group 1 and group 17 elements. -The elements are further grouped into 4 “blocks”. -Although positioned above Ne in group 18, helium is not considered a p-block element.

-Some commonly used group names:


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-The periodic table arranges elements according to the number of protons as well as the number and arrangement of electrons. -The following model treats the electrons as being in fixed shells, defined by the number n, around the nucleus. -The first shell (n = 1) can hold up to two electrons. -The second shell (n = 2) can hold up to eight electrons. -The third shell (n = 3) can hold up to eighteen electrons.

maximum # of electrons in shell n = 2n2 (where n = 1, 2, 3, …)

-Each shell is divided into n subshells. -The following table lists the names of the first four types of subshells and the max number of electrons in each:

Subshell Max # of Electrons in Subshell

s 2

p 6

d 10

f 14

-The first shell (n = 1) only has an s-type subshell, and consequently, can hold up to two electrons. -The second shell (n = 2) has both an s-type and p-type subshell, and therefore can hold up to eight electrons (two in the s subshell and six in the p subshell). -The third shell (n = 3) has s-type, p-type, and d-type subshells. -The fourth shell (n = 4) has four subshells, and so on.

Shell (n) Subshells Max # of Electrons in Subshells

1 s 2 2(1)2 = 2

2 s, p 2 + 6 = 8 2(2)2 = 8

3 s, p, d 2 + 6 + 10 = 18 2(3)2 = 18

4 s, p, d, f 2 + 6 + 10 + 14 = 32 2(4)2 = 32

…and so on


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-The subshells within a particular shell are labelled with the n value. ex. 3s, 3p, and 3d are the names given to the subshells when n = 3. Valence Electrons valence electrons → the number of electrons in the s and p subshells of the occupied electronic shell with the largest n for the neutral element ↘This shell is called the valence shell. -The electrons in the d and f subshells are not counted as valence electrons. -Valence electrons are important because they are the particles that are involved in the formation of bonds to other atoms. ex. In hydrogen and helium, the electrons are all in the n = 1 shell, therefore the n = 1 shell is the valence shell and the electrons in that shell (1 for hydrogen and 2 for helium) are valence electrons. ex. Oxygen has eight electrons, 2 in the n = 1 shell and 6 in the n = 2 shell.

-The n = 2 shell contains only 2s and 2p subshells, therefore, the electrons in the n = 2 shell are valence electrons and oxygen has 6 valence electron.

ex. O2- has 10 electrons, 2 in the n = 1 shell and 8 in the n = 2 shell. -n = 2 is the valence shell and O2- has 8 valence electrons ex. Mg2+ has 10 electrons, but the neutral Mg atom has 12 electrons.

-In the neutral Mg atom, 2 electrons are in the n = 1 shell, 8 in the n = 2 shell, and 2 in the n = 3 shell, therefore n = 3 is the valence shell and Mg has 2 valence electrons. -In the Mg2+ ion, the 2 valence electrons are taken away. -n = 3 is still the largest n for the neutral element, therefore n = 3 is still the valence shell and Mg2+ has 0 valence electrons.

-For neutral atoms of the s-block elements, the number of valence electrons is the same as the group #. -For neutral atoms of the p-block elements, the # of valence electrons is the group number – 10. -In ionic species, which have an unequal # of protons and electrons, electrons are added or removed from the valence shell of the neutral species. ex. O2- has 2 more electrons than neutral O which has 6 valence electrons, therefore the O2- has eight valence electrons, which is the max number of electrons for the n = 2 shell. ↘This is called a closed shell. -Elements can lose valence electrons to form cations with closed shells. ex. Mg can lose two valence electrons to form Mg2+ which has no electrons in its valence shell (n = 3). Effective Nuclear Charge (Zeff) -Many chemical species are positively charged, but in order for these cations to form, electrons must be removed from the valence shell of the neutral atom.


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-Since electrons are held to the nucleus by electrostatic forces of attraction, energy is required to move these electrons away. -Hydrogen only has one electron, therefore the only electrostatic forces present result from the attraction between the single electron and the single proton. -In lithium, the lone valence electron cannot “feel” the effect of the entire +3 charge in the nucleus because there is competition amongst three electrons. ↘There are also two electrons that are closer to the nucleus than the outermost electron. ↘These two electrons are called core electrons. ↘The core electrons “shield” the valence electron from “feeling” the full positive charge of the nucleus.

-All the elements in a group have the same Zeff. ↘Zeff is only used to understand properties from left to right on the periodic table. effective nuclear charge (Zeff) → the nuclear charge that the valence electron “feels” -An estimate of Zeff can be made with the following equation:

Zeff ≈ Z - S where Z = nuclear charge and S = # of core electrons

ex. The Zeff for a valence electron in lithium is 3 - 2 =1. -Zeff increases from left to right on the periodic table. Atomic Radius -The size of atoms is a factor when considering periodic trends. -Zeff can be used to rationalize trends in atomic radii. -The stronger the attractive force between the nucleus and the valence electrons, the smaller the atomic radius. -As Zeff increases, the electrons are pulled closer together towards the nucleus. -Therefore, moving from left to right within a period generally results in a reduction of atomic radius.


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-When going down a group, the number of shells containing electrons increases, therefore the atomic radius increases. -In cations, the fact that there are fewer electrons results in the electrons “feeling” more of the positive charge of the nucleus, therefore cations are smaller than the corresponding neutral atoms because they are pulled closer to the nucleus. -Anions have more electrons but the same # of protons, so the electrons “feel” less of the positive charge and are larger than the neutral atom. ex. S2- has a larger atomic radius than S because S2- has two more electrons and so the electrons are held less tightly because the # of protons stays the same. ↘The Zeff of both species is the same but spread out differently. ex. Na has a larger atomic radius than Na+ because in Na+, an electron is lost, and so the electrons are held more tightly towards the nucleus because the # of protons remains the same. ↘The Zeff is also different for these two species: Na → 11 - 10 = 1 Na+ → 11 - 2 = 9 Isoelectronic → same # of electrons isovalent → same # of valence electrons -In isoelectronic cations, the species with the largest Z will have the smallest ionic radius. -In isoelectronic anions, the species with the smallest Z will have the largest ionic radius.

Ionization Energy

ionization energy → the energy required to remove a single electron from an atom or ion in its gaseous state ↘Ionization energies are measured in the gas phase so that the energy due to interaction between species (in liquids or solids) does not affect the measurement. -Ionization energy is proportional to the magnitude of the electrostatic attraction between the electron being removed and the nucleus.


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-The electrostatic attraction is proportional to the nuclear charge felt by the electron (Coulomb’s Law). -Therefore, ionization energy should be directly proportional to Zeff and inversely proportional to the radius of the species. -The ionization energies generally increase from left to right across a row within the periodic table. -When the atomic radius of a species is smaller, it is harder to take away an electron. -Ionization energies generally decrease top to bottom within a particular group of the periodic table. ↘This can be rationalized by considering that the size of the atoms increases from top to bottom. ↘Electrons are more tightly held to the nucleus in elements higher up in the table. -This is because the energy of attraction between the nucleus and valence electrons is inversely proportional to the distance between them (Coulomb’s Law). Electron Affinity (EA) electron affinity → the energy change that results from the addition of the single electron to an atom or ion in its gaseous state ↘Electron affinities are measured in the gas phase so that the energy due to interaction between species (in liquids or solids) does not affect the measurement. -Generally, the addition of an electron to a neutral atom involves the release or energy. ↘This energy is directly proportional to Zeff. ↘The further on the right of the periodic table the element is, the higher the Zeff, the higher the EA. -Within a group, the EA decreases from top to bottom due to the weaker electrostatic attraction between the incoming electron and the nucleus as the atoms increase in size. -EA’s for neutral atoms are typically negative, due to the release of energy, however, when comparing EA’s for different species, only take into account the relative magnitudes. ex. The EA of chlorine is -349kJ/mol while bromine is -325kJ/mol. ↘Chlorine has a larger EA than bromine. -EA can be seen as the opposite of IE. -EA doesn’t always follow trends. Chemical Bonding -Most chemicals are formed from combinations of elements in defined proportions called compounds. compound → a charge neutral substance made up to two or more elements ex. NaCl, CH4, PCl3

-Atoms in a chemical compound are held together by bonds. -Chemical bonds consist of electrons that simultaneously feel the electrostatic forces of attraction from the positively charged nuclei of adjacent atoms. -Forming bonds in a compound helps the elements achieve a lower overall energy than they would in their free elemental form.


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-The three general types of chemical bonds are: ionic bond → oppositely charged ions are held together by electrostatic forces covalent bond → atoms are held together by the mutual attraction of a pair (or pairs) or electrons to the nuclei of adjacent atoms metallic bond → electrons are shared between many atoms simultaneously and are free to “flow” between the atoms ↘This free movement of electrons in a metal is the reason for the high electrical conductivity of metals. -Species possessing either an overall positive or negative charge are called ions. -Ions that have a positive charge are called cations. -Ions that have a negative charge are called anions. -Ions that contain a single atom are called simple ions. -Ions that contain multiple atoms are called complex or molecular ions. -A cation is not isolable in a condensed liquid or solid form unless there is an anion present to balance the charge. ex. Na+ will not be found in a bottle without a counterion such as Cl- which will make NaCl. Ionic Bonding -Consider the interaction of a lithium atom with a fluorine atom. -The lithium atom has one valence electron while the fluorine atom holds seven valence electrons. -When lithium transfers an electron to fluorine, the resultant species become isoelectronic with their nearest noble gas. -This transfer of electrons results in Li+ and F- ions which are attracted to each other by electrostatic forces or attraction in an ionic bond. -The empirical formula of this compound will be LiF. -For ionic compounds of the s- and p-block, the cation charge is the same as its group # and the anion charge is equal to 18 minus its group #. Covalent vs Ionic Bonding -In their compounds with p-block elements, the s-block elements typically form ionic bonds. -The reaction of sodium with chlorine results in a complete transfer of an electron from sodium to chlorine, giving sodium chloride which consists of Na+ and Cl- ions held together by electrostatic forces of attraction (represented by a dashed line):

Na+ ---- Cl- -In some cases, complete transfer of an electron will not happen between two elements. ex. H2 will not consist of H+ and H- ions because there is no reason for the two hydrogen atoms to have different numbers of valence electrons.


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-When this happens, a covalent bond is formed. -In covalent bonds, the electrons are attracted to both nuclei at the same time. -Since one electron comes from each atom to form a covalent bond, the electrons are said to “pair up”. -This “bond pair” is represented by a solid line:

H –– H

-In general, elements will form covalent bonds to achieve a closed shell. ex. Iodine will share one electron with a second iodine atom to form the molecule I2:

I –– I -This structure does not show all of the valence electrons and is therefore incomplete. -Represent the valence electron by drawing pairs of dots next to the elemental symbol:

-These electrons are not involved in bonding and are called lone pairs. -This diagram is called a Lewis diagram. -In some cases, an octet is achieved when more than one bond pair is present between atoms. ↘Two bond pairs is called a double bond. ↘Three pairs is called a triple bond. Polar Covalent Bonding -In homonuclear diatomic molecules (ex. H2, Cl2), the electrons are shared equally between the atoms, and the overall molecule is said to be nonpolar since the electrons are not polarized towards either atom of the molecule. -In heteronuclear diatomic molecules (ex. HF), the electrons are not shared equally, therefore one atom carries a partial positive charge and the other a partial negative charge. ↘This bond type is called polar covalent. -Partial positive charges are indicated with δ+. -Partial negative charges are indicated with δ-.

-In polar covalent bonds, electrons are polarized towards the atom with a greater electronegativity.


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electronegativity (chi - χ) → the power of an atom in a molecule to attract electrons to itself ↘Electronegativity is a molecular property and should not be confused with the atomic properties EA and IE. ↘This means you cannot discuss the electronegativity of an atom unless it is part of a molecule or molecular ion. -Electronegativity cannot be measured experimentally. bond dipole moment (µ) → the product of the charge (δ+ or δ-) on the two bonded atoms and the distance (d) between the two atoms

overall dipole moment → the vector sum of all the individual bond dipole moments in a molecule -Electronegativity values may be used to predict the polarity of a bond formed between any two s- or p- block elements. ex. The electronegativities of carbon and sulfur are the same (χ = 2.6). ↘It is expected that bonds between these atoms will be non-polar. ex. The electronegativity difference between hydrogen (χ = 2.2) and bromine (χ = 3.0) is relatively large. ↘One could predict that the bond in hydrogen bromide is a polar covalent bond. -The overall dipole moment (µ) is represented by placing a vector arrow along the direction of the bond dipole. -The arrowhead is on the δ- side and the crossed end is on the δ+ side. -When comparing different molecules, the length of the vector indicates the relative magnitudes of the dipole moment.

-The arrows indicate that HCl has a larger dipole moment than HBr. -There are two extremes of chemical bonding: purely covalent bonds (bonding electrons shared exactly equally between atoms as in homonuclear species such as [N2]

2- and O2) and purely ionic bonds (ie. one electron or more completely transferred between two atoms. -All heteronuclear compounds exhibit bonding between these two extremes. -In general, when the electronegativity difference between the elements involved in a bond is less than 1.6, a polar covalent bond will result. -If the difference is greater than 1.6, an ionic bond will result.

Bond Type Electronegativity Difference Bonding Electrons Examples


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Between Adjacent Atoms

Ionic ∆χ > 1.6 Completely transferred NaCl, MgO

Polar Covalent 0 < ∆χ ≤ 1.6 One from each HF, BrF

Non-Polar Covalent ∆χ = 0 One from each H2, F2

chapter 2: periodic properties: atoms and simple...

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