Gen chemistry periodic table trends
Which of the following is considered paramagnetic?
A paramagnetic electron is an unpaired electron. An atom is considered paramagnetic if even one orbital has a "net" spin. An atom could have ten diamagnetic electrons, but as long as it also has one paramagnetic electron, it is still considered a paramagnetic atom. •In normal triplet form, O2 molecules are paramagnetic. This means they form a magnet in the presence of a magnetic field because of the spin magnetic moments of the unpaired electrons in the molecule as well as the negative exchange energy between neighboring O2 molecules. Anytime two electrons share the same orbital, their spin quantum numbers have to be different. Whenever two electrons are paired together in an orbital, or their total spin is 0, and they are considered diamagnetic electrons. Atoms with all diamagnetic electrons are called diamagnetic atoms. Diamagnetic atoms repel magnetic fields. The unpaired electrons of paramagnetic atoms realign in response to external magnetic fields and are therefore attracted. Paramagnets do not retain magnetization in the absence of a magnetic field because thermal motion randomizes the spin orientations. •Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.
It has been observed that gaseous hydrogen is a very poor conductor of electricity, but a solution of hydrogen chloride in water is a good conductor. Which of the following reasons account for this?
Each of the hydrogen halides (Cl is a halogen) ionizes to at least some extent when it dissolves in water. H2O HCl(g)-------> H+(aq) + Cl-(aq) Ionization of hydrochloric acid: HCl(ℓ)+H2O(ℓ)→H3O+(aq)+Cl−(aq) HCl(g) is very soluble and is a very strong electrolyte. This means that it ionizes when dissolved in water, yielding aqueous ions which can conduct electricity. Halogen family consists of non-metallic elements. The halogen group (7) is the most electronegative in the periodic table, and all elements readily form halide ions X-. • Halogens are the only periodic table group that contains elements in all 3 familiar states of matter at standard temperature and pressure. • Halogens are electronegative and oxidizing elements, fluorine exceptionally so. They occur in nature as halides and form highly reactive diatomic molecules. • Molecular halides are formed with most nonmetals, ionic halides with metals. Some halides are good Lewis acids, and many halide complexes are known. • Halogens form acids when bonded to hydrogen. • Most halogens are typically produced from minerals or salts. • Halogens are toxic and middle halogens, chlorine, bromine, and iodine, are often used as disinfectants. Characteristics: • Halogens are highly reactive due to the high electronegativity, high effective nuclear charge. They can gain an electron by reacting with atoms of other elements. • Fluorine is one of the most reactive elements in existence, attacking otherwise inert materials such as glass and forming compounds with the heavier noble gases. • F2 > Cl2 > Br2 > I2 (oxidizing strength) • I- > Br- > Cl- > F- (reducing strength) • Halogens are poor conductors of heat and electricity.
Which of the following is the correct order of decreasing first ionization potential?
Effective nuclear charge = nuclear charge - shielding electrons. • Shielding electrons are those that stand between the nucleus and the electron we are interested in. • Shielding electrons are those that are in subshells closer to the nucleus (lower in energy) than the electron we are interested in. • The higher the effective nuclear charge for an electron, the more stable it is (higher ionization energy, not easily knocked off). • Effective nuclear charge increases for outer electrons as you go across (left to right) the periodic table. • Ionization potential (IP): The amount of energy required to take out the most loosely held electron from a neutral isolated atom. Factors affecting the IP: • Atomic size:-ionization potential is inversely proportional to atomic size. • Nuclear charge:-ionization potential is proportional to nuclear. • Shielding effect • Half filled subshells:- half filled subshells assume extra stability due to their symmetry. Hence it become more difficult to remove electrons from the half filled subshell and ionization potential become higher. IP generally increases this is due to a gradual increase in nuclear charge and decrease the atomic size of element. Elements of second period are Li, Be, B, C, N, O, F, Ne. So the order of first four elements should be: C>B>Be>Li. But this is not correct because first IP of boron is less than the first IP of Be. It is due to presence of one unpaired electron in p-subshell. Be:1s2 2s2 B:1s2 2s2 2p1
The two outside energy levels of an atom have the electronic structure s2 p6 d5 s2. To which family does the atom belong?
Electron configuration is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals. For example, the electron configuration of the neon atom is 1s2 2s2 2p6. Electronic configurations describe electrons as each moving independently in an orbital, in an average field created by all other orbitals. Standard notation is used to indicate the electron configurations of atoms and molecules. For atoms, the notation consists of a sequence of atomic orbital labels with the number of electrons assigned to each orbital placed as a superscript. • For example, hydrogen has one electron in the s-orbital of the first shell, so its configuration is written 1s1. Lithium has two electrons in the 1s-subshell and one in the (higher-energy) 2s-subshell, so its configuration is written 1s2 2s1. • Quantum numbers l, m, s, and number of quantum states (electrons) per orbital. • l is the angular momentum quantum number: l are integers that range from 0 to n-1. • spdf: l=0,1,2,3 for s,p,d,f respectively. • spdf designates subshells. • s subshells hold 1 orbital. p holds 3, d holds 5, f holds 7. • Each orbital holds a maximum of 2 electrons. • s subshells hold a maximum of 1x2=2 electrons, p: 3x2=6, d: 5x2=10, f: 7x2=14. • A generalized formula for the above pattern: for any subshell, 4l+2 electrons can be held. For a given shell, higher subshells have higher energy. • A low shell with a high subshell may be higher in energy than a higher shell with a low subshell. • m is the magnetic quantum number: m are integers that range from -l to l, including zero. • s is the spin quantum number: s is either +1/2 or -1/2. • Group 7, numbered by IUPAC nomenclature, is a group of elements in the periodic table. They are manganese (Mn), technetium (Tc), rhenium (Re), and bohrium (Bh). All known elements of group 7 are transition metals. • Like other groups, the members of this family show patterns in their electron configurations, especially the outermost shells resulting in trends in chemical behavior.
Which of the following elements demonstrates the highest second ionization energy?
Ionization energy is the energy needed to remove the most loosely bound electron from a neutral atom to form a positively charged ion. Alkali metals like Li, Na, K, and Rb have 1 valence electron that is easily removed. The next electron to be removed would be from a fully filled shell, requiring a very high second ionization energy. When comparing elements in the same group, the second ionization energy of an atom with a lower atomic number is considerably increased than the energy required to remove the second electron of an atom of a higher atomic number. K demonstrates a second ionazation energy higher than Rb. Valence electrons are electrons that are associated with an atom, and that can participate in the formation of a chemical bond. Alkaline earth metals like Be, Mg, Ca, and Sr have 2 electrons in the outermost shell. After removal of 1 electron from the valence shell, they all would only have 1 electron left, which can easily be removed in order to attain a stable electronic configuration. Therefore, the energy needed to remove the 2nd electron from the valence shell is less than the energy to remove the first. Example: It is more difficult to remove a second electron from K, because the electron would be in the lower energy 3p subshell compared to Rb, due to the 2nd electron would be from the 4p subshell.
Which of the following is the correct nomenclature for anion symbol N3-?
Ions are charged chemical species. Negatively charged ions are termed anions as they travel to the anode. Positively charged ions are called cations because they travel to the cathode of an electrolytic cell. • Common names for some ions are: Br- bromide I- iodide O2- oxide S2- sulfide N3- nitride • A salt is the solid which dissolves to produce these cations and anions. When a salt is dispersed in water (solvent) to form an aqueous solution of ions, also termed an ionic or electrolyte solution, there is an energetically favorable ion-dipole interaction.
Each of the following is a cation EXCEPT one. Which is the EXCEPTION?
Ions are charged chemical species. Positively charged ions are called cations because they travel to the cathode of an electrolytic cell. • Negatively charged ions are termed anions as they travel to the anode. • A salt is the solid which dissolves to produce these cations and anions. When a salt is dispersed in water (solvent) to form an aqueous solution of ions, also termed an ionic or electrolyte solution, there is an energetically favorable ion-dipole interaction. • Common names for some ions are: K+potassium ion Br-bromide Cs+cesium ion I-iodide Be2+beryllium ion O2-oxide Mg2+magnesium ion S2-sulfide Ca2+calcium ion N3-nitride Ba2+barium ion Al3+aluminum ion Ag+silver ion Zn2+zinc ion Cd2+cadmium ion
Each of the following is TRUE about the metals EXCEPT one. Which is the EXCEPTION?
Metals are solid materials (an element, compound, or alloy) that is typically hard, opaque, shiny, and features good electrical and thermal conductivity. Metals are generally malleable—that is, they can be hammered or pressed permanently out of shape without breaking or cracking—as well as fusible (able to be fused or melted) and ductile (able to be drawn out into a thin wire). 91 of the 118 elements in the periodic table are metals and are closely positioned to neighboring atoms in 1 of 2 common arrangements known as body-centered cubic and face-centered cubic. Metals are usually inclined to form cations through electron loss, reacting with oxygen in the air to form oxides. Non-metals are the chemical elements which mostly lack metallic attributes. Physical properties: • Are highly volatile • Have low elasticity • Are good insulators of heat and electricity Chemical properties: • Have high ionization energy • Have high electronegativity • Gain or share electrons when they react with other elements or compounds Seventeen elements are generally classified as non-metals: • Gases-hydrogen, helium, nitrogen, oxygen, fluorine, neon, chlorine, argon, krypton, xenon and radon • Liquid-bromine • Solids-carbon, phosphorus, sulfur, selenium, and iodine
Which of the following is the correct electron configuration for an atom of a noble gas like Argon (Ar)?
The electron configuration can be easily found when visualized on the periodic table. Argon, being a noble gas will have it's outermost shell full. Noble gases make a group of chemical elements with similar properties: (under standard conditions) They were once labeled group 0 in the periodic table because it was believed they had a valence of zero, meaning their atoms cannot combine with those of other elements to form compounds. The 6 noble gases that occur naturally are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive radon (Rn). Noble gases are typically highly unreactive except when under particular extreme conditions. As a result of a full shell, the noble gases can be written as a configuration with a full shell to form the noble gas notation. Noble gases have weak interatomic force, and consequently have very low melting and boiling points. They are all monatomic gases under standard conditions. Noble gases have the largest ionization potential among the elements of each period, which reflects the stability of their electron configuration and is related to their relative lack of chemical reactivity. Noble gases are dominated by the weak van der Waals forces between the atoms. Reactivity follows the order Ne < He < Ar < Kr < Xe < Rn
Among the alkali metals, the metal with the highest ionization potential is which of the following?
The ionization potential decreases from top to bottom in a given group with an increase in the size of the atom. Hence 'Li' possesses highest ionization potential among alkali metals. Alkali metals are a group in the periodic table consisting of the chemical elements lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). • This group lies in the s-block of the periodic table as all alkali metals have their outermost electron in an s-orbital. Properties: • They are all shiny, soft, highly reactive metals at standard temperatures and pressures. They readily lose their outermost electron to form cations with a +1 charge. • They can all be cut easily with a knife due to their softness, exposing a shiny surface that tarnishes rapidly in air due to oxidation. Because of their high reactivity, they must be stored under oil to prevent reaction with air and are found naturally only in salts and never as the free element. • In modern IUPAC nomenclature, the alkali metals comprise the group 1 elements, excluding hydrogen (H), which is nominally a group 1 element but not normally considered to be an alkali metal. • All of the alkali metals react with water, with the heavier alkali metals reacting more vigorously than the lighter ones.
Each of the following is TRUE about the chemical characteristics of alkaline earth metals EXCEPT one. Which statement is the EXCEPTION?
This group lies in the s-block of the periodic table as all alkaline earth metals have their outermost electron in an s-orbital. GROUP 2: ALKALINE EARTH metals. They are all shiny, silvery-white, somewhat reactive metals at STP and readily lose their two outermost electrons to form cations with charge 2+ and an oxidation state, or oxidation number of +2. These are Be, Mg, Ca, Sr, Ba, and Ra. Chemical characteristics: • Have relatively low densities, melting points, and boiling points. • React with the halogens to form the alkaline earth metal halides. • All are ionic crystalline compounds (except for beryllium chloride, which is covalent). • All except beryllium react with water to form strongly alkaline hydroxides. • The alkaline metals have the second-lowest first ionization energies in their respective periods of the periodic table because of their somewhat low effective nuclear charges and the ability to attain a full outer shell configuration by losing just two electrons. • The second ionization energy of alkaline metals are also somewhat low. • Beryllium is an exception: it does not react with water or steam, and its halides are covalent. • All the alkaline earth metals have two electrons in their valence shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions.
As the oxidation state of Mn increases from +2 to +7, acidic character of the oxide will do which of the following?
Transition elements form groups 3-11 in the d-block. They have distinct chemical characteristics resulting from the progressive filling of the d shells. These include the occurrence of variable oxidation states, and compounds with structures and physical properties resulting from partially filled d orbitals. Vertical trends: Elements of the 3d series are chemically very different from those in the 4d and 5d series, showing weaker metallic and covalent bonding, stronger oxidizing properties in high oxidation states, and the occurrence of many more compounds with unpaired electrons. Horizontal trends: Electropositive character declines towards the right of each series. Elements become less reactive and their compounds show a tendency Towards 'softer' behavior. Later elements in the 4d and 5d series are relatively more inert. Electron configurations of the neutral atoms are complex and have both d and s electrons in outer shells. For example, in the 3d series most atoms have the configuration (3d)n(4s)2, where n increases from one to 10; chromium and copper are exceptions with (3d)5(4s)1and (3d)10(4s)1, respectively. The configurations depend on a balance of 2 factors: • 3d orbitals are progressively stabilized relative to 4s across the series; • Repulsion between electrons is large in the small 3d orbitals, and so minimum energy in the neutral atom is achieved in spite of (i) by putting one or 2 electrons in the 4s orbitals. Characteristic properties: Properties shared by the transition elements that are NOT found in other elements, which results from the partially filled d shell are: • The formation of compounds whose color is due to d-d electronic transitions • The formation of compounds in many oxidation states, due to the relatively low reactivity of unpaired d electrons • The formation of many paramagnetic compounds due to the presence of unpaired d electrons. A few compounds of main group elements are also paramagnetic (e.g. nitric oxide, oxygen). • The transition elements have low ionization energies. They exhibit a wide range of oxidation states or positively charged forms. The positive oxidation states allow transition elements to form many different ionic and partially ionic. • Although Mn+2 is the most stable ion for manganese, the d-orbital can be made to remove 0 to 7 electrons. The increase in the number of oxidation states from Sc to Mn. All possible states exhibited by only Mn.
