Ch7 periodic properties of elements Flashcards
mendeleev3
- and meyer noticed similar chem and phys properties recur periodically when the elements are arranged in order of atomic weight. Their tables were the forerunners of the modern PT
- Mendeleev insisted that elements with similar characteristics be listed in the same family, so there were many blank spaces in his table (elements that were unknown at the time)
- Mendeleev was able to predict the existence and properties of unknown elements (predicted gallium and called it eka-aluminum (under-aluminum))
moseley
determined that each element produces x rays of a unique frequency, and the frequency increased as atomic mass increased. He assigned a unique whole number (atomic number) to each element to arrange the frequencies. The atomic number was equal to the number of protons/electrons in the atom.
effective nuclear charge5
- Zeff, is the net electric field as if it results from a single positive charge located at the nucleus.
- Zeff is smaller than the actual nuclear charge because Zeff also account for the repulsion of the electron by the other electrons in the atom
- greater Zeff=greater attraction between orbital and nucleus=lower energy for orbital (ns
shield
Core electrons partially cancel the attraction of the valence electron and the nucleus=they partially shield/screen the outer electron from the attraction of the nucleus.
nonbonding radii
The closest distances separating the nuclei during the collisions determine the apparent radii/nonbonding atomic radius/van der Waals radii of the atoms
bonding radii
- distances separating nuclei of atoms when they are chemically bonded together. Shorter than nonbonding radius.
- length of bonds is equal to the sum of the the radii of the two atoms involved
atomic radius2
- Atomic radius increases going down a group, because principal quantum number n increases. Outer electrons have a greater probability of being farther from nucleus, so atom increases in size
- Atomic radius decreases going across a period because Zeff increases, so force of attraction increases and valence electrons are drawn closer to the nucleus
ionic radii2
- Cations are smaller than parent atoms because there are less electrons. Anions are bigger than their parent atoms because there are more electrons.
- For ions with the same charge, size increases as you go down a column in PT because n increases so the orbital increases
isoelectronic series
group of ions all containing same number of electrons (O2-, F-, Na+, Mg2+ all have 10 electrons). If isoelectronic ions are listed in increasing atomic number, (therefore increasing Zeff), the ionic radius decreases because electrons are more strongly attracted to the nucleus
ionization energy5
- minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. Greater ionization energy=more difficult to remove electron
- First ionization energy: I1, energy needed to remove first electron from a neutral atom (Na–>Na++e-).
- I1
ionization E special cases3
- Representative elements (s and p block) show larger range of values of I1 than transition metals. Transition metals increase slowly across the PT
- Decrease in Ionization energy from Be to B because the third valence electron of B must occupy the 2p subshell, which is at a higher energy than 2s.
- Decrease in ionization energy from N to O because repulsion of paired electrons in the p^4 configuration (two of the electrons have to share an orbital, so there is repulsion)
adding/removing an electron2
- Electrons removed from an atom to form a cation are removed first from the occupied orbitals having largest quantum number n. Fe ([Ar]3d64s2)–>Fe2+([Ar]3d6). Electrons are removed from 4s orbitals before 3d orbitals, even though electron configurations says 4s is filled before 3d.
- Electrons added to an atom to form an anion are added to the empty/partially filled orbital with lowest value of n F(1s22s22p5)–>F-(1s22s22p6)
electron affinity2
- energy change when an electron is added to a gaseous atom. Measures attraction/affinity of the atom for the added electron. Greater attraction between atom and added electron=more negative atom’s electron affinity
- usually increase left to right(become more negative) until noble gases. Electron affinities do not change greatly going down a period because electron is being added to a subshell with same number of electrons in it as element above.
ionization E vs electron affinity
Ionization energy:Cl–>Cl++e- E=1251kJ/mol
Electron affinity: Cl+e- –>Cl- E=-349kJ/mol
Ionization energy measure the ease with which an atom loses an electron, electron affinity measure ease with which an atom gains an electron
electron affinities special cases3
- can have a positive value for some elements like noble gases and Be and Mg, because the added electron would reside in a previously empty subshell that is higher in energy and would take a lot of energy=unfavorable
- Halogens have most negative electron affinities because they only need one electron to have a full p shell
- Electron affinities in group 5A (N,P, As, Sb): all have half filled p orbitals so adding an electron would put it in an already occupied orbital=larger repulsion. These elements have positive electron affinities (N) or less negative than neighbors to the left
metallic character
the more an element exhibits the physical and chemical properties of metal. Increases down a column and decreases across a row
metal properties5
- shiny luster, various colors, most are silvery.
- Solids are malleable (pounded into thin sheets) and ductile (drawn into wires).
- Good conductors of heat/electricity.
- Most metal oxides are ionic solids that are basic.
- Tend to form cations in aqueous solutions. (low Ionization energies)
metal oxides are basic2 equations
metal oxide+water=metal hydroxide
Metal oxide+acid=salt +water
nonmetal properties5
- no luster, various colors. (vary in appearance)
- Brittle, some are hard, some are soft.
- Poor conductors of heat/electricity.
- Most nonmetal oxides are molecular substances that form acidic solutions.
- Tend to form anions or oxyanions in aqueous solution. (high electron affinities)
nonmetal oxides are acidic equations 2
nonmetal oxide+water=acid.
Nonmetal oxide+base =salt+water
group 1A properties4
- soft metallic solids, with silver luster and
- high conductivity.
- They have low I1 ionization energies, so they are very reactive.
- They emit a characteristic color when placed in a flame
group 1A equations3
- Alkali metals react with hydrogen to form hydrides and sulfur to form sulfides 2M(s)+H2(g)→ 2MH(s)
- They react with water to produce hydrogen gas and a solution of metal hydroxide 2M(s)+2H2O(l)–> 2MOH(aq)+H2(g). These reactions are exothermic
- When reacting with oxygen, they either form metal oxides(O2- ion, ex: Li), peroxides (O22-ion, ex:Na), or superoxides (O2-ion, ex:K,Rb, Cs).
group 2A properties4
- harder, denser and
- melt at higher temp than alkali metals.
- Their I1 energies are not as low as alkali metals, so they are not as reactive.
- They also give off colors when in a flame
group 2A equations2
- react with steam to form metal oxide + hydrogen gas: Mg(s) + H2O(g) → MgO(s)+H2(g).
- They react with water to form metal hydroxide and hydrogen gas: Ca(s)+2H2O(l) → Ca(OH)2(aq)+H2(g). These equations show the alkaline earth metal tendency to lose two outer s electrons and form 2+ ions
oxygen things2
- two allotropes: O2(common) and O3(ozone). To go from O2 to O3 is endothermic rxn
- Compounds of peroxide and superoxide react with themselves to produce an oxide and O2 in an exothermic rxn: 2H2O2(aq)–>2H2O(l)+O2(g)
sulfur things2
- most common allotrope: S8 which is usually written as just S(s).
- Sulfur reacts with O2 to form SO2(sulfur dioxide)
group 7A things4
- halogens are all nonmetals and all diatomic (F2, Cl2..).
- At room temp, F and Cl are gas, Br is liquid, I is solid.
- They have highly negative electron affinities, so they tend to gain electrons from other elements to form halide ions: X-
- react with H2 to form gaseous hydrogen halide compounds: H2(g)+X2–>2HX(g)
cl reacts with water
Cl2(g)+H2O(l)–> HCl(aq)+HOCl(aq)
group 8A things3
- noble gases and are all nonmetal gases.
- They have completely full s and p subshells so their I1 ionization energies are very high, so they are unreactive.
- Used to be called inert gases.
