ch.7 chem – Flashcards

+3

+2

+1

+3

were known last

were known secondly

were know first

published similar works on arranging the elements by order of increasing atomic weight

element Cn( 112; and not included in some periodic tables) are the newest elements

discover atomic numbers Bombarding different elements with high-energy electrons, Moseley found that each element produced X-rays of a unique frequency and that the frequency generally increased as the atomic mass increased. He arranged the X-ray frequencies in order by assigning a unique whole number, called an atomic number, to each element.

tells us that the strength of the interaction between two electrical charges ( for ex, the protons and electrons) depends on the magnitudes of the charges ( nuclear charge) and on the distance between them.

is the positive charge created by the nucleus, and that attracts electrons

is smaller than the actual nuclear charge ( Zeff < Z ) because the effective nuclear charge includes the effect of the other electrons in the atom.

Zeff= Z- S where Z= number of protons S= positive screening constant;represents the portion of the nuclear charge that is screened from a valence electron by the other electrons in the atom; usually is substituted by number of core electrons

but they do affect the value of S slightly; this is why we use Slater's rule, because its equation includes the part of the total nuclear charge that is also affected by the valence electrons

the general trend in orbital energies (ns < np < nd) in a many-electron atoms. The greater attraction between the 2s electron and the nucleus leads to a lower energy for the 2s orbital than for the 2p orbital.

INCREASES FROM LEFT TO RIGHT across any period of the periodic table. Although the number of core electrons stays the same across the period, the number of protons increases. The valence electrons added to counterbalance the increasing nuclear charge screen one another ineffectively. Thus, Zeff increases steadily - You can always remember this, by Looking at Li and Be... Li has an Zeff= 3-2= +1 while Be has an Zeff=4-2=+2 INCREASES SLIGHTLY AS WE GO DOWN A COLUMN because the more diffuse core electron cloud is less able to screen the valence electrons from the nuclear charge.

we can use his approach if we limit ourselves to elements that do not have electrons in d or f subshells. - Electrons for which the principal quantum number n is larger than the value of n for the electron of interest contribute 0 to the value of S. - Electrons with the same value of n as the electron of interest contribute 0.35 to the value of S. - Electrons for which n is 1 less than n for the electron of interest contribute 0.85, while those with even smaller values of n contribute 1.00. Ex; For a v. e- in fluorine, Slater's rules tells us that S= (0.35 x6) +(0.85 x 2) = 3.8 (Slater's rules ignore the contribution of an electron to itself in screening; therefore, we consider only six n=2 electrons, not all seven). Thus, Zeff = Z-S = 9-3.8= 5.2 +

exactly replicate the values of Zeff obtained from more sophisticated calculations, both methods effectively capture the periodic variation in Zeff. So, don't freak out if you don't get the same Zeff as in the book!

Imagine a collection of argon atoms in the gas phase. When two of these atoms collide with each other, they ricochet apart like colliding billiard balls. This ricocheting happens because the electron clouds of the colliding atoms cannot penetrate each other to any significant extent. The shortest distance separating the two nuclei during such collisions is twice the radii of the atoms

based on the distance separating the nuclei when two atoms are bonded to each other

for any atom in a molecule is equal to half of the nucleus-to-nucleus distance d.

the bonding atomic radius

(2.66 angstrom/2)= 1.33 angstrom

For example, the Cl-- Cl bond length in Cl2 is 1.99 Å, so a bonding atomic radius of 0.99 Å is assigned to Cl. In CCl4 the measured length of the C-Cl bond is 1.77 Å, very close to the sum (0.77 + 0.99Å) of the bonding atomic radii of C and Cl. P.S- C has a bonding atomic radius of 0.77Å

The angstrom ( 1 Å=10^-10 m) is a convenient metric unit for atomic measurements of length. IT IS NOT AN IS UNIT. The most commonly used SI unit for atomic measurements is the picometer (1 pm = 10^-12 m; 1 Å = 100 pm).

Within each group, bonding atomic radius tends to increase from top to bottom. This trend results primarily from the increase in the principal quantum number (n) of the outer electrons. As we go down a column, the outer electrons have a greater probability of being farther from the nucleus, causing the atomic radius to increase. Within each period, bonding atomic radius tends to decrease from left to right. The major factor influencing this trend is the increase in effective nuclear charge Zeff across a period. The increasing effective nuclear charge steadily draws the valence electrons closer to the nucleus, causing the bonding atomic radius to decrease.

For ions carrying the same charge, ionic radius increases as we move down a column in the periodic table. Ionic radius have the same trend as atomic radius, but the ions have to have the same charge!

a group of ions all containing the same number of electrons. For example, each ion in the isoelectronic series O^-2, F^- , Na^+, Mg^+2, Al^+3 has 10 electrons.In any isoelectronic series we can list the members in order of increasing atomic number; therefore, nuclear charge increases as we move through the series. Ionic radius decreases with increasing nuclear charge as the electrons are more strongly attracted to the nucleus ----------------------> Increasing nuclear charge O^-2 F^- Na^+ Mg^+2 Al^+3 ------- Decreasing atomic Radius---->

with increasing nuclear charge as the electrons are more strongly attracted to the nucleus.

This trend exists because with each successive removal, an electron is being pulled away from an increasingly more positive ion, requiring increasingly more energy.

Zeff nuclear charge, increases across a period and Boron goes before Carbon, which means Carbon has a greater Zeff. If Carbon had a greater Zeff before, now would be even greater, since 6 prontons exerting force over its four remaining electrons. If you notice, also B^+ and C^+2, the products of the ionization energies, have the same number of electrons, which mean they are an isoelectronic series. --- increasing nuclear charge ( atomic number)---> B^+ C^+2 --- Decreasing ionic radius---> Meaning harder time to pull the other electrons from the tightly bound to the nucleus Carbon

the harder is to remove an electron

the easier to remove an electron

I1 generally increases as we move across a period; it is harder to remove a noble gas electron than al alkali metal electron I1 generally decreases as we move down any column in the periodic table; it is easier to remove an elctron of an atom whole electrons are farther away from the nucleus. The s- and p-block elements show a larger range of I1 values than do the transition metal elements. Generally, the ionization energies of the transition metals increase slowly from left to right in a period. The f-block metals also show only a small variation in the values of I1.

The decrease in Ionization energy ( easiness to remove an electron) from Be [ Ne] 2s^2 to B [ Ne] 2s^2 2p^1 is because B would need to get rid of its 1 electron to be more stable.

is a measure of the energy change associated with removing an electron from the atom to form a cation.

energy is released when an electron is added For example, the addition of an electron to a chlorine atom is accompanied by an energy change of -349 Kj/mol , the negative sign indicating that energy is released during the process.We therefore say that the electron affinity of Cl is -349 KJ/mol

Ionization energy measures the ease with which an atom loses an electron, whereas electron affinity measures the ease with which an atom gains an electron.

THE MORE NEGATIVE THE ATOM'S ELECTRON AFFINITY. For some elements, such as the noble gases, the electron affinity has a positive value, meaning that the anion is higher in energy than are the separated atom and electron:

THE HIGHER THE AFFINITY

means that an electron will not attach itself to an Ar atom; the ion is unstable and does not form.

The halogens, which are one electron shy of a filled p subshell, have the most negative electron affinities. By gaining an electron, a halogen atom forms a stable anion that has a noble-gas configuration.

the electrons are placed in a higher energy level far from the nucleus, thus a decrease from its pull.

a more positive ionization energy means a higher energy to remove an electron

can be pounded into thin sheets

(of a metal) able to be drawn out into a thin wire.

are ionic solids that are basic

are molecular substances that form acidic solutions

and therefore tend to form cations relatively easily. As a result, metals are oxidized (lose electrons) when they undergo chemical reactions.

to lose electrons

is the best indicator of whether an element behaves as a metal or a nonmetal.

boron (B) silicon( Si), germanium ( Ge), arsenic (As), antimony (Sb), tellurium(Te), polonium (Po)

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+2

a binary compound of a halogen with another element or group.

Those that dissolve in water react to form metal hydroxides, as in the following example Metal oxide + water --->metal hydroxide Na2O(s) + H2O(l)---> 2 NaOH(aq)

O^2-(aq) + H2O(l)---> 2 OH-(aq)

Metal oxide + acid----> salt + water NiO (s) + 2 HNO3 ( aq) ---> Ni (NO3)2 (aq) + H20 (l)

H2 , N2, O2 , F2, Cl2 , Br2 ( liquid) , I2 ( volatile solid)

are typically molecular substances that tend to be gases, liquids, or low-melting solids at room temperature.

Nonmetal oxide + water ---> acid CO2 (g) + H2O (l) ---> H2CO3( aq) P4O10 (s) + 6H2O (l) ---> 4H3PO4 (aq)

Nonmetal oxide + base---> salt + water CO2(g) + 2 NaOH(aq) ---> Na2CO3(aq) + H2O(l)

have low densities and melting points exist in nature only as compounds.

hydrogen gas and a solution of an alkali metal hydroxide: 2 M(s) + 2 H2O(l) ----> 2 MOH(aq) + H2(g) M representing any alkali metal

4 Li (s) + O2 (g) ----> 2 Li2O (s) Lithium Oxide

For example Na forms sodium peroxide 2 Na (s) + O2 (g) ----> Na2O2 sodium peroxide

K (s) + O2 (g) ----> KO2 (s) potassium superoxide

the alkaline earth metals are harder and more dense and melt at higher temperatures.

- Hydrogen ionization energy is more comparable to that of the halogens than of the alkali metals. Instead, hydrogen shares its electron with nonmetals and thereby forms molecular compounds. - hydrogen reacts with active metals to form solid metal hydrides that contain the hydride ion, H^-1. The fact that hydrogen can gain an electron further illustrates that it is not truly an alkali metal. - In addition to its ability to form covalent bonds and metal hydrides, probably the most important characteristic of hydrogen is its ability to lose its electron to form a cation.

O2 ( dioxygen) and O3 (ozone)

different forms of the same element in the same state.

the product is less stable than the reactant

2 H2O2(aq) ----> 2 H2O(l) + O2(g) DELTA H° = -196.1 kJ ENDOTHERMIC

Even though solid sulfur consists of S8 rings, we usually write it simply as S(s) in chemical equations to simplify the stoichiometric coefficients.

down the column: H2O > H2S > H2Se > H2Te, with H2O, water, being the most stable of the series.

increase with increasing atomic number.