Chemistry Interim: Flashcards

Dalton (solid sphere) → Thomson (plum pudding) → Rutherford (nucleus) → Bohr (energy levels) → quantum mechanical (electron cloud).

Most went straight through: mostly empty space. A few were deflected: positive charge is concentrated. Very few bounced back: the nucleus is tiny, dense and positive.

Orbit (Bohr): a fixed circular path. Orbital (quantum): a region where an electron is probably found.

Thomson. Thomson's cathode ray tube (1897): the ray bent toward a positive plate, so it was made of negative particles.

The gold foil (alpha scattering) experiment. Rutherford, 1911. A few alpha particles bounced back, so something tiny, dense and positive had to be in the way.

plum pudding model. Electrons embedded in a sphere of positive charge, like fruit in a pudding.

The positive charge and most of the mass are packed into a tiny nucleus. Only a head-on hit with something small, dense and positive could turn an alpha particle around.

most of the atom is empty space. They hit nothing. The nucleus is a tiny target.

The charge of an electron. It gave the charge of one electron, which with Thomson's result also gave the electron's mass.

It cannot explain why orbiting electrons do not lose energy and spiral into the nucleus. It also could not explain line spectra. Bohr's fixed energy levels answered both.

The line spectrum of hydrogen. Only certain colors appear, so only certain energy changes are possible, so the levels must be fixed.

It only works for atoms with one electron, such as hydrogen. It fails for the spectra of larger atoms, and electrons do not really follow neat circular paths.

Chadwick. Chadwick, 1932, by bombarding beryllium with alpha particles.

a region where an electron is likely to be found. An orbital is a probability region (electron cloud). A fixed circular path is a Bohr orbit.

Atoms cannot be divided. Atoms are made of protons, neutrons and electrons. Isotopes also show that atoms of one element are not all identical.

Dalton, Thomson, Rutherford, Bohr, quantum. Ball, pudding, nucleus, levels, cloud.

Proton: +1, 1 amu, nucleus. Neutron: 0, 1 amu, nucleus. Electron: −1, about 1/1836 amu, electron cloud.

Protons = atomic number. Neutrons = mass number − atomic number. Electrons = protons − charge.

The proton. The number of protons is the atomic number. Change it and the element changes.

Protons and neutrons. Electrons are outside the nucleus, in the electron cloud.

0. Neutral. Proton +1, electron −1.

The electron. An electron has about 1/1836 the mass of a proton.

protons + neutrons. Electrons are too light to count. Neutrons = mass number − atomic number.

In the nucleus. Protons and neutrons are in the nucleus, and each is about 1836 times heavier than an electron.

tiny, but it holds nearly all the mass. The nucleus is about 100 000 times smaller than the atom. Most of the atom is empty space.

nucleons. Both live in the nucleus.

Isotopes: same protons, different neutrons. Ions: same protons, different electrons.

Positive ion: element name + "ion" (sodium ion). Negative ion: ending changes to -ide (chloride ion).

neutrons. Same protons (same element), different neutrons, so a different mass number.

electrons. The nucleus does not change. Only the electron count does.

¹²₆C and ¹⁴₆C. Same protons (6), different neutrons (6 and 8). The second pair are different elements; the last two are an atom and its ion.

a cation. Cation = positive (lost electrons). Anion = negative (gained electrons).

Chloride ion. A single-atom negative ion ends in -ide.

Sodium ion. A positive ion keeps the element's name. Sodium has only one charge, so no Roman numeral.

2−. Electrons are negative, so gaining two makes the charge 2−.

They have the same number of protons and electrons. Chemistry is decided by the electrons. Neutrons only change the mass.

Sum of (isotope mass × fractional abundance). Fraction = percent ÷ 100, and the fractions add to 1.

Let x and (1 − x) be the fractions. mass₁·x + mass₂·(1 − x) = average. Solve for x.

the weighted average of the masses of the element's natural isotopes. Each isotope counts in proportion to its abundance.

Ne-20. The average is always closest to the most abundant isotope.

divide it by 100 to get a decimal fraction. 75.78% becomes 0.7578.

1 − x. The fractions must add up to 1.

It is a weighted average of isotopes with different masses. No single chlorine atom has a mass of 35.45. It is the average over the natural mixture.

Fusion: light nuclei join (Sun, stars); needs extreme temperature and pressure. Fission: a heavy nucleus splits (power plants). Both release huge energy and form new elements.

The mass numbers (top) add to the same total on both sides, and so do the atomic numbers (bottom).

two light nuclei combining to form a heavier nucleus. Fusion joins. Fission splits.

In the Sun and other stars. Stars fuse hydrogen into helium. Power plants use fission.

Nuclei are positive and repel, so they must move very fast to get close enough to fuse. Like charges repel. Only very fast (very hot) nuclei collide hard enough to join.

helium. Hydrogen fuses to helium and releases the energy that makes the Sun shine.

The mass numbers and the atomic numbers. Add the top numbers and the bottom numbers separately on each side.

A small amount of mass is converted into energy. The products have slightly less mass than the reactants. E = mc².

The nucleus changes, so a new element forms. Chemical reactions only rearrange electrons, and the elements stay the same.

a heavy nucleus splitting into smaller nuclei. For example, uranium-235 splitting when a neutron hits it.

far more energy. Millions of times more energy for the same mass of fuel.

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p. Capacities: s 2, p 6, d 10, f 14.

Aufbau: lowest energy first. Pauli: at most 2 electrons per orbital, opposite spins. Hund: one electron in each orbital of a sublevel before pairing.

Count the electrons to find the element, then compare with the filling order. An electron in a higher sublevel while a lower one is not full means excited.

Put the previous noble gas (end of the row above) in brackets, then write only the electrons after it. [He] 2, [Ne] 10, [Ar] 18, [Kr] 36.

2. Two electrons with opposite spins (Pauli exclusion principle).

6. Three orbitals × 2 electrons. s holds 2, d holds 10, f holds 14.

10. Five orbitals × 2 electrons.

4s. 4s is lower in energy than 3d. There is no 3f.

each orbital in a sublevel gets one electron before any orbital gets two. One in each box first, all with the same spin. Then pair.

an orbital holds at most two electrons, with opposite spins. One up arrow and one down arrow per box, never two the same.

electrons fill the lowest-energy sublevel first. Build up from 1s, following the diagonal rule.

[Ar] 4s¹ 3d⁵. An exception: a half-filled 3d sublevel is extra stable. Copper is the other one: [Ar] 4s¹ 3d¹⁰.

the same number of electrons but more energy. An electron has moved to a higher level. Nothing was added or removed.

Excitation: the electron absorbs energy and jumps up. De-excitation: it falls back and emits a photon with energy equal to the gap between the levels.

E = hf and c = λf. Higher energy means higher frequency and shorter wavelength. Violet is high energy; red is low.

It jumps to a higher energy level. This is excitation. The atom is now in an excited state.

When an electron falls from a higher level to a lower level. This is de-excitation. The energy leaves as a photon.

the difference in energy between the two levels. E(photon) = E(high) − E(low).

Violet. Violet has the shortest wavelength and the highest frequency.

increases. c = λf and c is constant, so they change in opposite directions.

increases. E = hf.

the ground state. Any higher arrangement is an excited state.

a shorter wavelength. More energy means higher frequency and shorter wavelength.

Energy levels are fixed, so only certain energy changes are possible. Quantized levels give only certain photon energies, so only certain colors.

excited electrons fall back to lower energy levels. The flame excites the electrons. The color comes out on the way back down.

Red. Red is about 700 nm, violet about 400 nm.

unstable: the electron soon falls back and emits light. Atoms return to the ground state almost at once.

An element is present only if every one of its lines appears in the unknown. Leftover lines belong to another element.

410 nm (violet), 434 nm (blue-violet), 486 nm (blue-green), 656 nm (red). All are drops to n = 2.

Each element has its own unique set of lines. Different elements have different energy levels, so different lines.

When all of the element's lines appear in the unknown's spectrum. Every line must match. One match is not enough.

Hydrogen is present, with at least one other element. The extra lines must come from something else.

all the colors with no gaps. White light through a prism gives a continuous spectrum.

a glowing (excited) gas of an element. Excited atoms emit only their own particular colors.

one particular electron transition (energy drop). Each drop between two levels gives photons of one energy, so one line.

Their energy levels are different. Different nuclear charge and electron arrangement give different level spacings.

n = 2. Drops to n = 1 are ultraviolet and drops to n = 3 are infrared.

They match the lines in its spectrum with those of known elements. The same matching rule you use on the exam.

Decreases across a period (more protons, same shielding). Increases down a group (more energy levels, more shielding).

Cation smaller than its atom; anion larger. For ions with the same number of electrons, more protons means smaller.

decreases. More protons pull the same energy level in closer.

increases. Each row adds an energy level.

More protons pull the electrons in the same energy level closer. Nuclear charge rises while shielding stays about the same.

Each row adds an energy level, so the outer electrons are farther away and more shielded. Distance and shielding outweigh the extra protons.

smaller. It has lost electrons, often a whole energy level, and the same protons pull on fewer electrons.

larger. Extra electrons repel each other, and the same protons hold more electrons.

It has more protons pulling on the same number of electrons. 11 protons against 9, for the same 10 electrons.

Inner electrons reducing the pull of the nucleus on the outer electrons. Core electrons sit between the nucleus and the valence electrons.

Both increase across a period and decrease down a group. F is the most electronegative; He has the highest ionization energy.

Claim → cause (protons, shielding or energy levels) → effect (how strongly the outer electrons are held).

The jump comes when a core electron is removed, so the number of electrons before the jump is the number of valence electrons.

remove the outermost electron from a gaseous atom. The more tightly the electron is held, the higher the ionization energy.

how strongly an atom attracts shared electrons in a bond. It only applies to atoms in a bond.

increases. Greater nuclear charge with the same shielding holds the electrons more tightly.

decreases. The outer electron is farther away and more shielded.

F. Top right of the table (ignoring the noble gases).

The outer electron is farther from the nucleus and more shielded, so it is held less tightly. Distance and shielding win over the larger nuclear charge.

Nuclear charge increases while shielding stays about the same. The outer electrons feel a stronger pull, so more energy is needed to remove one.

a high ionization energy and a high electronegativity. Small atom, tight hold on electrons.

low ionization energies, so they form cations. They lose electrons easily.

He. Helium is the smallest atom with a full level and no shielding.

an ionic bond. Above about 1.7 the electron is transferred. This is the link between Unit 2 and Unit 3.

Less than 0.5: nonpolar covalent. 0.5 to 1.7: polar covalent. More than 1.7: ionic. Metal + metal is metallic.

Metal first: ionic, no prefixes, Roman numeral if the metal has more than one charge, polyatomic ions keep their names. Two non-metals: prefixes and -ide.

When it has polar bonds and the dipoles do not cancel: a lone pair on the central atom, or outer atoms that differ.

Ionic. A metal (Na) with a non-metal (Cl): ΔEN = 3.0 − 0.9 = 2.1, so electrons are transferred and the ions attract.

Metallic. Copper is a metal: its + ions sit in a sea of free-moving electrons.

Nonpolar covalent. Two identical non-metal atoms: ΔEN = 0, so the pair is shared equally.

Polar covalent. Two non-metals with ΔEN between 0.5 and 1.7: shared unequally, pulled toward O.

They are transferred from the metal to the non-metal. The metal loses electrons and becomes a + ion; the non-metal gains them and becomes a − ion. The opposite charges attract.

They move freely among the positive metal ions. This "sea of electrons" explains why metals conduct and can be bent.

Nonpolar covalent. Below 0.5 counts as nonpolar covalent.

Ionic. Above about 1.7 between a metal and a non-metal, the electron is transferred: ionic.

Metallic bonds. Two metals: an alloy, held by metallic bonding.

Cl, because it is more electronegative. The shared pair is pulled toward the more electronegative atom, Cl (3.0 vs 2.1).

Both ionic and covalent. Ionic between Na⁺ and NO₃⁻; covalent between N and O inside the nitrate ion.

Whether the shared electrons are shared unequally or equally. Both share electrons. In a polar bond one atom pulls the pair closer.

The difference in electronegativity between the two bonded atoms. ΔEN = larger EN − smaller EN. Under 0.5 nonpolar, 0.5–1.7 polar covalent, over 1.7 ionic.

One [Ca]²⁺ and two [Cl]⁻, each Cl with 8 dots. Ca gives away 2 electrons, one to each Cl. Charges: 2+ and 2 × 1− add to zero.

24. C 4 + O 3 × 6 = 22, plus 2 for the 2− charge = 24.

8. N 5 + H 4 × 1 = 9, minus 1 for the + charge = 8.

C. Carbon forms the most bonds (4) and is the least electronegative. H is never central.

It can form only one bond. H needs only 2 electrons, so it makes one bond and is always on the outside.

2. 8 electrons: 2 O–H bonds use 4, the other 4 are 2 lone pairs on O.

A triple bond. 10 electrons: N≡N with one lone pair on each N gives both atoms 8.

4. Each dash is one shared pair = 2 electrons. A double bond is 2 pairs.

Square brackets with the charge outside. The charge belongs to the whole ion, so it goes outside the brackets, top right.

The double bond can go to any of the three O atoms; the real ion is a blend. That is resonance. All three N–O bonds are identical in the real ion.

+1. FC = valence − lone-pair electrons − bonds = 5 − 0 − 4 = +1.

C≡C. Higher bond order = more shared pairs = shorter and stronger.

Move a lone pair from an outer atom to make a double bond. You never add electrons. You turn a lone pair into a shared pair.

Metallic. Only metals have electrons that are free to move in the solid.

The ions are free to move once it melts. Ions carry the current. In the solid they are locked in the lattice.

A shifted layer puts like charges next to each other, and they repel. A knock moves one layer by one ion: + meets + and − meets −, so the crystal splits.

The ion layers slide while the electron sea keeps holding them. Malleable and ductile: nothing has to break when the layers move.

The forces between its molecules are weak. Melting separates molecules from each other. It does not break the bonds inside them.

Ionic. High melting point, brittle, conducts only when the ions can move: ionic.

Metallic. Lustre, malleability and conduction as a solid all come from the electron sea.

Molecular covalent. Small molecules with weak forces between them: low melting and boiling points.

Covalent network. Every atom is joined to its neighbors by covalent bonds, as in diamond.

The bond gets shorter and stronger. More shared pairs pull the two nuclei closer and take more energy to break.

Water molecules are polar, so they attract each other strongly. Polar molecules attract each other; nonpolar ones have only weak attractions.

A repeating lattice of + and − ions. There is no "NaCl molecule": each Na⁺ is surrounded by Cl⁻ and each Cl⁻ by Na⁺.

Dinitrogen pentoxide. Two non-metals: prefixes. The "a" of penta- is dropped before "oxide".

Carbon monoxide. No mono- on the first element, but the second always takes a prefix.

Magnesium chloride. Metal + non-metal: no prefixes. Mg only forms 2+, so no Roman numeral either.

Iron(III) chloride. Three Cl⁻ make 3−, so the iron is Fe³⁺. Iron has more than one charge, so the numeral is needed.

Copper(II) oxide. O is 2−, so the single copper ion must be Cu²⁺.

Copper(I) oxide. O is 2−, shared by two copper ions, so each is Cu⁺.

Sodium nitrate. NO₃⁻ is the polyatomic ion nitrate. Use its own name.

Calcium carbonate. CO₃²⁻ is carbonate. Calcium only forms 2+, so no numeral.

Ammonium sulfate. Two polyatomic ions: ammonium NH₄⁺ and sulfate SO₄²⁻. No prefixes in ionic names.

Phosphorus trichloride. Two non-metals: prefixes, with no mono- on the first element.

Sulfur hexafluoride. Two non-metals: hexa- for 6.

Al₂O₃. Al³⁺ and O²⁻: criss-cross the charge numbers to get Al₂O₃ (6+ and 6−).

Ca(OH)₂. Ca²⁺ needs two OH⁻. Two of the same polyatomic ion go in parentheses.

FeSO₄. Fe²⁺ and SO₄²⁻ balance one to one.

N₂O₄. Read the prefixes as subscripts: di = 2, tetra = 4.

When both elements are non-metals (a covalent compound). Ionic names never use prefixes: the charges fix the ratio.

When the metal can form more than one charge. Iron, copper and most transition metals. Not Na, Ca, Al, Zn or Ag.

Each iron ion has a 3+ charge. The numeral is the charge of the metal ion. The formula is Fe₂O₃.

Sulfite. -ite has one oxygen fewer than -ate: sulfate is SO₄²⁻. Sulfide is S²⁻.

Hydrogen sulfide. When hydrogen comes first it takes no prefix.

Cu(NO₃)₂. Cu²⁺ needs two NO₃⁻, in parentheses.

Copper: its free electrons carry current and it can be drawn into wire. Metallic bonding gives conduction in the solid and ductility.

A plastic (molecular covalent): it insulates and bends. Molecular covalent materials have no free charges, so they insulate, and they are flexible.

An ionic ceramic: very high melting point. Strong attraction between the ions means it survives very high temperatures.

Diamond (covalent network): the hardest. In a covalent network every atom is locked in by covalent bonds.

Atoms of a different size stop the layers sliding easily. Sliding layers are what make a pure metal soft. Disrupting them makes it harder and stronger.

Copper and zinc. Brass = Cu + Zn. Bronze = Cu + Sn.

Copper and tin. Bronze = Cu + Sn. Brass = Cu + Zn.

Chromium. Stainless steel is iron + carbon + chromium + nickel.

It is brittle and would shatter. Ionic materials are hard but brittle: fine for a tile, wrong for anything that takes impacts.

They are poor conductors of heat. The metal body conducts heat to the food; the covalent handle keeps the heat from your hand.

A nonpolar covalent material. Water is polar. Nonpolar materials such as wax or nonstick coatings do not attract it.

A homogeneous mixture of a metal with other elements. No new compound forms: the atoms are mixed, still held by metallic bonding.

Ionic. Dissolving frees the ions, and moving ions carry the current.

They conduct heat well. The free electrons carry heat as well as charge.

Covalent network, such as quartz (SiO₂). Network solids are hard with very high melting points, and many are transparent.

List the properties the job needs. Then match those properties to the bond type that gives them.

2. Two double bonds, and each counts as one domain.

Tetrahedral. 4 bonded atoms and no lone pairs on C.

Trigonal pyramidal. 4 domains: 3 atoms + 1 lone pair.

Bent. 4 domains: 2 atoms + 2 lone pairs.

Linear. 2 domains and no lone pairs on C.

Trigonal planar. 3 domains and no lone pair on B.

Bent. 3 domains: 2 atoms + 1 lone pair on S.

109.5°. Four domains point to the corners of a tetrahedron.

About 104.5°. Two lone pairs squeeze the angle below 109.5°.

120°. Three domains spread evenly in a plane.

Tetrahedral. Electron geometry counts all 4 domains, lone pair included.

Tetrahedral. 2 bonded atoms + 2 lone pairs = 4 domains.

Bent, a little less than 120°. Only 2 atoms are attached, so the atoms form a bent shape.

The lone pair repels more than a bonding pair. Lone pairs take more space and push the bonds closer together.

The bond comes out of the page, toward you. Plain = in the page, wedge = toward you, dash = away.

Trigonal planar. N has 3 bonded O atoms and no lone pair: 3 domains.

Trigonal pyramidal. 26 electrons: P has 3 bonds + 1 lone pair = 4 domains.

It is linear, so the two equal dipoles cancel. Two equal arrows pointing in opposite directions add to zero.

It is bent, so the two O–H dipoles add instead of cancelling. The shape decides: the two arrows point partly the same way.

No: ΔEN is 0.4. A bond with ΔEN below 0.5 counts as nonpolar.

O. O is more electronegative (3.5 vs 2.1), so it pulls the shared pair.

Its bonds are polar but the molecule is nonpolar. C–Cl is polar (ΔEN 0.5). Four equal dipoles in a tetrahedron cancel.

Its bonds are nonpolar and the molecule is nonpolar. C–H has ΔEN 0.4: no polar bonds, so nothing to add up.

Yes: the outer atoms are not all the same. The C–F dipoles are stronger than the C–Cl dipoles, so they cannot cancel.

Yes: the two ends are different. Linear cancels only when both outer atoms are the same, as in CO₂.

The more electronegative atom. That is where the shared electrons are pulled: the δ− end.

Polar. The lone pair makes the shape unsymmetric, so the dipoles do not cancel.

NF₃: it has a lone pair and is trigonal pyramidal. BF₃ is trigonal planar and symmetric, so its three dipoles cancel.

Yes: it is bent. S has a lone pair, so SO₂ is bent and its S–O dipoles add.

NH₃, because it is polar like water. "Like dissolves like": polar dissolves in polar.

Its shape: whether the bond dipoles cancel. Molecular polarity is the vector sum of the bond dipoles.