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Bagrut Chemistry practice questions

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Question 1Structure of the Atom and the Periodic Table

Which statement correctly describes the quantum numbers of an electron in a 3p orbital?

  1. n = 3, l = 0, and the magnetic quantum number mℓ is always 0
  2. n = 3, l = 3, and the magnetic quantum number mℓ may be −3 to +3
  3. n = 3, l = 2, and the magnetic quantum number mℓ may be −2, −1, 0, +1, or +2
  4. n = 3, l = 1, and the magnetic quantum number mℓ may be −1, 0, or +1 Correct answer.

Explanation

The letter p corresponds to l = 1, and the principal quantum number n is the leading numeral, so n = 3. For l = 1 the allowed values of mℓ are −1, 0, and +1, giving three p orbitals.

Question 2Structure of the Atom and the Periodic Table

How many electrons can occupy the entire 4f subshell?

  1. 14 Correct answer.
  2. 2
  3. 6
  4. 10

Explanation

An f subshell has l = 3, so there are 2l + 1 = 7 orbitals, and each orbital holds two electrons with opposite spins. The capacity is therefore 7 × 2 = 14 electrons.

Question 3Structure of the Atom and the Periodic Table

Which ground-state electron configuration is correct for a neutral chlorine atom (Z = 17)?

  1. 1s² 2s² 2p⁶ 3s² 3p⁴ 4s¹
  2. 1s² 2s² 2p⁶ 3s² 3p⁶
  3. 1s² 2s² 2p⁵ 3s² 3p⁶
  4. 1s² 2s² 2p⁶ 3s² 3p⁵ Correct answer.

Explanation

Filling 17 electrons by the Aufbau principle gives 1s² 2s² 2p⁶ 3s² 3p⁵, ending one electron short of the argon core. This matches chlorine's position in group 17 with seven valence electrons.

Question 4Structure of the Atom and the Periodic Table

Across a period from left to right, which trend is observed for atomic radius?

  1. It decreases because the number of shells increases
  2. It decreases because effective nuclear charge increases while electrons enter the same shell Correct answer.
  3. It stays constant because the number of protons equals the number of electrons
  4. It increases because more electrons are added to the atom

Explanation

Protons are added across a period while electrons enter the same principal shell, so the effective nuclear charge rises and pulls the electron cloud inward. The result is a smaller atomic radius from left to right.

Question 5Structure of the Atom and the Periodic Table

How many neutrons are present in an atom of ³⁷Cl?

  1. 54
  2. 17
  3. 20 Correct answer.
  4. 37

Explanation

The superscript 37 is the mass number, which equals protons plus neutrons, and chlorine's atomic number gives 17 protons. Subtracting gives 37 − 17 = 20 neutrons.

Question 6Structure of the Atom and the Periodic Table

Which element has the ground-state configuration [Ar] 4s² 3d¹⁰ 4p³?

  1. Arsenic (Z = 33) Correct answer.
  2. Antimony (Z = 51)
  3. Phosphorus (Z = 15)
  4. Selenium (Z = 34)

Explanation

The [Ar] core accounts for 18 electrons, and the remaining 15 electrons (4s² 3d¹⁰ 4p³) bring the total to 33. Element 33 is arsenic, which lies in group 15 below phosphorus.

Question 7Structure of the Atom and the Periodic Table

Which of the following is the correct set of values for the spin quantum number ms of a single electron?

  1. 0 to n − 1
  2. +1 or −1
  3. −l to +l
  4. +½ or −½ Correct answer.

Explanation

The spin quantum number describes the intrinsic angular momentum of an electron and takes only the two values +½ and −½. These two values allow each orbital to hold at most two electrons.

Question 8Structure of the Atom and the Periodic Table

Which element has the highest first ionisation energy?

  1. Neon Correct answer.
  2. Potassium
  3. Sodium
  4. Argon

Explanation

First ionisation energy rises across a period and falls down a group. Neon and argon both have full shells, but neon's valence electrons are in the n = 2 shell, closer to the nucleus, so neon has the highest value.

Question 9Structure of the Atom and the Periodic Table

An element has the successive ionisation energies (kJ/mol): 590, 1150, 4940, 6480. In which group of the periodic table is it most likely found?

  1. Group 1
  2. Group 2 Correct answer.
  3. Group 13
  4. Group 14

Explanation

The third ionisation energy jumps sharply compared with the second, showing that the third electron is removed from a filled inner shell. Two relatively easy removals indicate two valence electrons, which is characteristic of group 2.

Question 10Structure of the Atom and the Periodic Table

Which electron configuration represents a chromium atom (Z = 24) in its ground state?

  1. [Ar] 4s¹ 3d⁵ Correct answer.
  2. [Ar] 4s² 3d⁴
  3. [Ar] 4s⁰ 3d⁶
  4. [Ar] 4s² 3d⁶

Explanation

Chromium is an exception to the strict Aufbau order because a half-filled d subshell with 4s¹ 3d⁵ has extra exchange stabilisation and lower energy. Its actual ground-state configuration is therefore [Ar] 4s¹ 3d⁵.

Question 11Structure of the Atom and the Periodic Table

Which particle is isoelectronic with the K⁺ ion (Z = 19)?

  1. Ar
  2. Na⁺
  3. Cl⁻ Correct answer.
  4. Ca²⁺

Explanation

A potassium atom has 19 electrons, so K⁺ has 18 electrons. A chlorine atom has 17 electrons, and adding one gives Cl⁻ a total of 18, making it isoelectronic with K⁺.

Question 12Structure of the Atom and the Periodic Table

Which statement explains why the first ionisation energy of magnesium is higher than that of aluminium?

  1. Aluminium has a half-filled 3p subshell that is unusually stable
  2. Aluminium has a larger nuclear charge that is fully shielded by its inner electrons
  3. Magnesium has a smaller number of protons than aluminium, so it holds electrons more tightly
  4. Magnesium's outermost electron is in a filled 3s orbital, whereas aluminium's is in a higher-energy 3p orbital Correct answer.

Explanation

Removing an electron from magnesium means breaking into a stable, filled 3s subshell, while aluminium loses a 3p electron that is higher in energy and better shielded. This makes aluminium's first ionisation energy lower despite its larger nuclear charge.

Question 13Structure of the Atom and the Periodic Table

Which set of quantum numbers is forbidden for an electron in an atom?

  1. n = 2, l = 1, mℓ = −1
  2. n = 2, l = 2, mℓ = 0 Correct answer.
  3. n = 3, l = 2, mℓ = 2
  4. n = 4, l = 0, mℓ = 0

Explanation

The orbital quantum number l may take only values from 0 to n − 1, so for n = 2 the maximum l is 1. The combination n = 2, l = 2 therefore cannot exist.

Question 14Structure of the Atom and the Periodic Table

Which group of the periodic table contains elements whose atoms have a general valence configuration of ns² np⁵?

  1. Group 15
  2. Group 17 Correct answer.
  3. Group 18
  4. Group 16

Explanation

A valence configuration of ns² np⁵ has seven valence electrons, one short of a full octet. This is characteristic of the halogens in group 17, which readily gain one electron.

Question 15Structure of the Atom and the Periodic Table

Which property generally increases down group 1 of the periodic table?

  1. First ionisation energy
  2. Effective nuclear charge felt by the valence electron
  3. Atomic radius Correct answer.
  4. Electronegativity

Explanation

Each step down group 1 adds a new principal shell, so the valence electron lies farther from the nucleus and the atomic radius increases. The added shielding also lowers ionisation energy and electronegativity.

Question 16Structure of the Atom and the Periodic Table

An atom's emission spectrum shows a line at a wavelength corresponding to an energy difference of 3.0 × 10⁻¹⁹ J. Which process produces this line?

  1. An electron falling from a higher energy level to a lower one Correct answer.
  2. A proton in the nucleus changing its energy state
  3. An electron moving from a lower energy level to a higher one
  4. An electron being removed completely from the atom

Explanation

Emission lines arise when an excited electron drops to a lower energy level and releases the energy difference as a photon of definite frequency. Absorption lines, by contrast, correspond to upward transitions.

Question 17Structure of the Atom and the Periodic Table

Which statement about the shapes of atomic orbitals is correct?

  1. Both s and p orbitals are spherically symmetric around the nucleus
  2. A p orbital consists of four lobes arranged tetrahedrally
  3. An s orbital is spherically symmetric, while a p orbital has two lobes separated by a nodal plane Correct answer.
  4. An s orbital has two lobes, while a p orbital is spherically symmetric

Explanation

The s orbital has no angular node and is spherically symmetric about the nucleus. Each p orbital has one angular node, giving two lobes of opposite sign along one axis.

Question 18Structure of the Atom and the Periodic Table

Which element would be expected to have the largest atomic radius?

  1. Beryllium
  2. Caesium Correct answer.
  3. Strontium
  4. Barium

Explanation

All four elements are in group 2 or the adjacent group 1, and radius increases down a group. Caesium lies in period 6, the lowest of these positions, so its valence electron occupies the n = 6 shell and gives the largest radius.

Question 19Structure of the Atom and the Periodic Table

Which explanation best accounts for the very high third ionisation energy of sodium compared with its first two?

  1. Sodium has only two electrons in total, so a third cannot be removed
  2. The third electron is removed from the 3p orbital, which is unusually stable
  3. The third ionisation energy is high because sodium becomes a cation
  4. The first two electrons come from the 3s orbital, but the third requires breaking into the stable neon core Correct answer.

Explanation

Sodium's configuration is [Ne] 3s¹, so the first ionisation removes the single 3s electron and the second removes an electron from the stable neon core. The large jump reflects the much greater energy needed to disrupt a full shell.

Question 20Structure of the Atom and the Periodic Table

Which pair of elements has atoms with the same number of unpaired electrons in their ground state?

  1. Fluorine and neon
  2. Boron and carbon
  3. Carbon and oxygen Correct answer.
  4. Nitrogen and oxygen

Explanation

Carbon (2p²) has two unpaired p electrons by Hund's rule, and oxygen (2p⁴) also has two unpaired electrons. Both therefore have two unpaired electrons despite their different electron counts.

Question 21Structure of the Atom and the Periodic Table

Which isoelectronic series is arranged in order of increasing ionic radius?

  1. F⁻ < Na⁺ < Mg²⁺ < Al³⁺
  2. Al³⁺ < Mg²⁺ < Na⁺ < F⁻ Correct answer.
  3. Mg²⁺ < Al³⁺ < F⁻ < Na⁺
  4. Na⁺ < F⁻ < Al³⁺ < Mg²⁺

Explanation

All four ions have ten electrons, so the nuclear charge determines the size: more protons pull the same electron cloud in more tightly. Nuclear charges are Al (13), Mg (12), Na (11), and F (9), so radius increases in the order Al³⁺ < Mg²⁺ < Na⁺ < F⁻.

Question 22Structure of the Atom and the Periodic Table

A neutral atom in its ground state has exactly three electrons with n = 3 and l = 2. Which element could it be?

  1. An element with atomic number between 13 and 18
  2. An element with atomic number between 31 and 36
  3. An element with atomic number between 19 and 22
  4. An element with atomic number between 23 and 30 Correct answer.

Explanation

Electrons with n = 3 and l = 2 are 3d electrons, and having exactly three of them means the configuration ends in 3d³. The 4s subshell fills first, so this corresponds to vanadium (Z = 23) and the range of elements near it in the first transition series.

Question 23Structure of the Atom and the Periodic Table

Which observation provides the strongest evidence that electrons occupy discrete energy levels rather than a continuum?

  1. The mass of an atom is concentrated in a small nucleus
  2. A gas discharge tube emits light only at certain sharp wavelengths Correct answer.
  3. A heated metal filament glows with a continuous range of colours
  4. The electrical conductivity of a metal increases as it is cooled

Explanation

Sharp line spectra mean photons are emitted only at specific energies, which requires the atom's energy to take only discrete values. A continuous spectrum would be expected if electron energies formed a continuum.

Question 24Structure of the Atom and the Periodic Table

Which statement correctly compares the effective nuclear charge felt by a 3s electron in magnesium with that felt by a 3s electron in sodium?

  1. It is the same in both because both electrons occupy a 3s orbital
  2. It is greater in sodium because sodium has fewer protons pulling on the electron
  3. It is greater in magnesium because the extra proton is added to the same shell without adding a new inner shell Correct answer.
  4. It is greater in magnesium because magnesium has a larger atomic radius

Explanation

Going from sodium to magnesium adds one proton and one 3s electron in the same shell, so the added electron shields only partially. The net effective nuclear charge on the 3s electrons therefore increases, which also explains magnesium's smaller radius.

Question 25Chemical Bonding and Intermolecular Forces

Which type of bonding is best described as the electrostatic attraction between a lattice of positive ions and a sea of delocalized electrons?

  1. Ionic bonding
  2. Covalent bonding
  3. Hydrogen bonding
  4. Metallic bonding Correct answer.

Explanation

Metallic bonding arises when metal atoms release valence electrons into a shared delocalized pool, and the resulting attraction between the fixed cations and mobile electrons holds the lattice together. This delocalization also explains the electrical and thermal conductivity of metals.

Question 26Chemical Bonding and Intermolecular Forces

A bond forms between potassium (electronegativity 0.8) and chlorine (electronegativity 3.0). Based on the electronegativity difference, the bond is best classified as:

  1. Ionic Correct answer.
  2. Metallic
  3. Purely covalent
  4. Polar covalent

Explanation

The electronegativity difference is 2.2, which exceeds the usual threshold of about 1.7, so electron transfer dominates and the interaction is treated as ionic. The large difference means chlorine attracts the shared electron density so strongly that discrete K+ and Cl- ions form.

Question 27Chemical Bonding and Intermolecular Forces

Which intermolecular force is present in all molecular substances, including nonpolar ones such as methane?

  1. Hydrogen bonding
  2. Ion-dipole interactions
  3. London dispersion forces Correct answer.
  4. Dipole-dipole interactions

Explanation

London dispersion forces arise from instantaneous fluctuations in electron distribution that create temporary dipoles, so they operate in every molecule regardless of polarity. In methane they are the only intermolecular force available.

Question 28Chemical Bonding and Intermolecular Forces

Which of the following species has a trigonal planar molecular geometry according to VSEPR theory?

  1. NH3
  2. BF3 Correct answer.
  3. H2O
  4. CH4

Explanation

BF3 has three bonding pairs and no lone pairs on boron, so the three electron domains adopt a 120-degree trigonal planar arrangement. VSEPR geometry depends on the number of electron domains and how many are lone pairs.

Question 29Chemical Bonding and Intermolecular Forces

Which property is characteristic of ionic compounds in the solid state?

  1. They are brittle and shatter under stress. Correct answer.
  2. They have low melting points.
  3. They conduct electricity well.
  4. They dissolve readily in nonpolar solvents.

Explanation

In an ionic solid, shifting layers by stress brings ions of like charge into alignment, causing repulsion and cleavage, so the crystal is brittle. The strong electrostatic lattice also gives high melting points and poor conductivity while ions are locked in place.

Question 30Chemical Bonding and Intermolecular Forces

Which molecule is polar overall?

  1. CO2
  2. CCl4
  3. H2O Correct answer.
  4. BF3

Explanation

Water has a bent geometry, so its two polar O-H bond dipoles do not cancel and the molecule has a net dipole moment. Symmetrical molecules with identical polar bonds can have zero net dipole.

Question 31Chemical Bonding and Intermolecular Forces

Which statement about a coordinate (dative) covalent bond is correct?

  1. The bond involves transfer of electrons between ions.
  2. Both shared electrons come from the same atom. Correct answer.
  3. Each atom contributes one electron to the bond.
  4. The bond exists only between identical atoms.

Explanation

In a coordinate bond, one atom donates a lone pair that is shared with an electron-deficient partner, as in the formation of the ammonium ion from NH3 and H+. Once formed, the bond is equivalent to an ordinary covalent bond.

Question 32Chemical Bonding and Intermolecular Forces

Which factor best explains why the boiling point of HF is much higher than that of HCl?

  1. HF molecules form hydrogen bonds with one another. Correct answer.
  2. HF is a stronger acid than HCl.
  3. HF has a smaller bond length than HCl.
  4. HF has a larger molar mass than HCl.

Explanation

Fluorine is highly electronegative and small, so the H-F bond allows strong hydrogen bonding between molecules, adding substantial energy to the boiling process. HCl, with chlorine too large and less electronegative, shows only dipole-dipole and dispersion forces.

Question 33Chemical Bonding and Intermolecular Forces

In the molecule HCN, what is the hybridization of the carbon atom?

  1. sp2
  2. sp3d
  3. sp Correct answer.
  4. sp3

Explanation

Carbon in HCN forms a single bond to H and a triple bond to N, giving two electron domains, which corresponds to sp hybridization and a linear geometry. The number of sigma-bonding domains plus lone pairs determines the hybridization.

Question 34Chemical Bonding and Intermolecular Forces

A student compares the melting points of NaCl and NaBr. Which prediction and reasoning are correct?

  1. NaBr melts higher because it has a larger lattice energy.
  2. NaCl melts higher because Cl- is smaller and its lattice energy is greater. Correct answer.
  3. Both melt at the same temperature because both are 1:1 salts.
  4. NaBr melts higher because Br- is more polarizable.

Explanation

Lattice energy increases as ion size decreases, because smaller ions allow closer approach and stronger electrostatic attraction. Chloride is smaller than bromide, so NaCl has the greater lattice energy and the higher melting point.

Question 35Chemical Bonding and Intermolecular Forces

Which molecule has the largest London dispersion forces among the following?

  1. C2H6
  2. C3H8
  3. C4H10 Correct answer.
  4. CH4

Explanation

Dispersion forces increase with the number of electrons and the surface area available for contact, both of which grow with molecular size. Butane, the largest of these alkanes, therefore has the strongest dispersion forces and the highest boiling point.

Question 36Chemical Bonding and Intermolecular Forces

Which species has a bond angle closest to 109.5 degrees?

  1. CH4 Correct answer.
  2. NH3
  3. CO2
  4. H2O

Explanation

Methane has four identical bonding pairs with no lone pairs, so the tetrahedral angle is the ideal 109.5 degrees. Lone pairs compress bond angles in ammonia and water, while CO2 is linear at 180 degrees.

Question 37Chemical Bonding and Intermolecular Forces

Why does magnesium oxide have a much higher melting point than sodium chloride?

  1. MgO involves 2+ and 2- ions, giving much stronger electrostatic attraction. Correct answer.
  2. MgO is held together by covalent bonds.
  3. MgO has weaker ionic bonds because Mg2+ is small.
  4. MgO has larger ions, increasing lattice energy.

Explanation

Lattice energy depends on the product of ionic charges, so a 2+ with a 2- ion produces roughly four times the electrostatic attraction compared with a 1+ with a 1- ion of similar size. The stronger lattice in MgO requires far more energy to break.

Question 38Chemical Bonding and Intermolecular Forces

Which intermolecular force is primarily responsible for the solubility of ethanol in water?

  1. London dispersion forces between hydrocarbon chains
  2. Ion-dipole interactions
  3. Hydrogen bonding between ethanol and water molecules Correct answer.
  4. Metallic bonding

Explanation

Ethanol's hydroxyl group can both donate and accept hydrogen bonds with water, and this favorable interaction compensates for the energy needed to separate the pure liquids. The short hydrocarbon chain is small enough not to disrupt this solvation.

Question 39Chemical Bonding and Intermolecular Forces

According to molecular orbital theory, why is the O2 molecule paramagnetic?

  1. It has a bond order of one.
  2. It contains unpaired electrons in pi antibonding orbitals. Correct answer.
  3. It contains an odd number of electrons.
  4. It has a net dipole moment.

Explanation

Filling the molecular orbitals of O2 places two electrons singly in the degenerate pi antibonding orbitals, giving two unpaired electrons and a paramagnetic ground state. Lewis structures fail to predict this behavior, which MO theory explains.

Question 40Chemical Bonding and Intermolecular Forces

Which pair of molecules is held together in the liquid state primarily by dipole-dipole forces rather than hydrogen bonds?

  1. HF and NH3
  2. HCl and H2S Correct answer.
  3. CH3OH and H2O
  4. NH3 and H2O

Explanation

Hydrogen bonding requires H directly bonded to N, O, or F. In HCl and H2S the hydrogen is bonded to chlorine or sulfur, so the dominant intermolecular attractions are dipole-dipole and dispersion forces.

Question 41Chemical Bonding and Intermolecular Forces

Which statement correctly describes the relationship between bond order, bond length, and bond strength?

  1. Higher bond order gives longer, weaker bonds.
  2. Higher bond order gives shorter but weaker bonds.
  3. Bond order does not affect bond length.
  4. Higher bond order gives shorter, stronger bonds. Correct answer.

Explanation

Increasing bond order means more shared electron pairs between the nuclei, which pulls the atoms closer and requires more energy to break the bond. Thus a triple bond is shorter and stronger than a double bond, which in turn exceeds a single bond.

Question 42Chemical Bonding and Intermolecular Forces

In a sample of solid argon, which intermolecular force holds the atoms together?

  1. London dispersion forces Correct answer.
  2. Hydrogen bonding
  3. Covalent bonding
  4. Ionic attraction

Explanation

Argon atoms are monatomic and nonpolar, so the only attraction between them is the weak, transient dispersion force arising from fluctuating electron clouds. This is why argon solidifies only at very low temperatures.

Question 43Chemical Bonding and Intermolecular Forces

Which of the following correctly ranks the species by increasing carbon-oxygen bond length?

  1. CO < CO3^2- < CO2
  2. CO3^2- < CO2 < CO
  3. CO < CO2 < CO3^2- Correct answer.
  4. CO2 < CO < CO3^2-

Explanation

Bond length decreases as bond order increases. Carbon monoxide has a triple bond (order about 3), carbon dioxide has double bonds (order 2), and the carbonate ion has resonance-averaged bonds of order about 1.33, giving the longest bonds.

Question 44Chemical Bonding and Intermolecular Forces

Which statement best explains why the boiling point of water is anomalously high compared with H2S?

  1. Water is a linear molecule.
  2. Each water molecule can form up to four hydrogen bonds due to its two donors and two acceptors. Correct answer.
  3. Water has weaker dispersion forces than H2S.
  4. Water has a higher molar mass than H2S.

Explanation

Oxygen's high electronegativity and two lone pairs let each water molecule donate two O-H hydrogens and accept two hydrogen bonds, creating an extensive network. Sulfur is less electronegative and larger, so H2S shows only weak dipole-dipole and dispersion forces.

Question 45Chemical Bonding and Intermolecular Forces

The carbonate ion, CO3^2-, is described by three equivalent resonance structures. What does this imply about its carbon-oxygen bonds?

  1. The bonds alternate rapidly between single and double.
  2. All three C-O bonds are double bonds.
  3. One C-O bond is a double bond and two are single bonds, permanently.
  4. All three C-O bonds are identical with bond order about 1.33. Correct answer.

Explanation

Resonance means the true structure is a weighted average of the contributing structures, so the three bonds are equivalent with a bond order between one and two. The delocalized pi electrons are spread over all three oxygen atoms.

Question 46Chemical Bonding and Intermolecular Forces

Which observation is best explained by the presence of delocalized electrons in a metal?

  1. Metals conduct electricity in the solid state and are malleable. Correct answer.
  2. Metals form brittle crystals.
  3. Metals have high melting points.
  4. Metals dissolve readily in water.

Explanation

Delocalized electrons are free to move under an applied potential, giving electrical conductivity, and they allow layers of cations to slide without breaking directed bonds, giving malleability. These properties follow directly from the electron-sea model.

Question 47Chemical Bonding and Intermolecular Forces

Two compounds have similar molar masses: propan-1-ol (CH3CH2CH2OH) and ethyl methyl ether (CH3CH2OCH3). Propan-1-ol boils much higher. Which explanation is correct?

  1. Propan-1-ol molecules hydrogen bond through the O-H group, while the ether has no O-H bond. Correct answer.
  2. The ether has a much larger molar mass.
  3. The ether cannot experience dispersion forces.
  4. Propan-1-ol is ionic and the ether is covalent.

Explanation

Hydrogen bonding requires hydrogen covalently bonded to an electronegative N, O, or F atom. Propan-1-ol has an O-H group that can donate hydrogen bonds, while the ether's oxygen has lone pairs but no attached hydrogen, so it can only accept, not donate.

Question 48Chemical Bonding and Intermolecular Forces

The bond angle in H2O is about 104.5 degrees while that in NH3 is about 107 degrees. Which explanation is most consistent with VSEPR theory?

  1. Oxygen has two lone pairs, which repel bonding pairs more strongly than the single lone pair on nitrogen. Correct answer.
  2. The O-H bond is longer than the N-H bond.
  3. Water has a linear electron-domain geometry.
  4. Nitrogen is more electronegative than oxygen.

Explanation

Lone pairs occupy more space than bonding pairs, so each additional lone pair pushes the bonding pairs closer together. Oxygen has two lone pairs versus one on nitrogen, so water's angle is compressed more than ammonia's.

Question 49Stoichiometry and the Mole Concept

כמה מולים של אטומי חמצן מכילים 1.5 מול של מולקולות O₂?

  1. 3.0 מול Correct answer.
  2. 6.0 מול
  3. 0.75 מול
  4. 1.5 מול

Explanation

כל מולקולת O₂ מכילה שני אטומי חמצן, ולכן מספר המולים של אטומים הוא כפול ממספר המולים של מולקולות. מכפילים 1.5 מול מולקולות ב-2 ומקבלים 3.0 מול אטומי חמצן. אין צורך במסת המולרית בחישוב זה, משום שהמעבר הוא בין מולקולות לאטומים בלבד.

Question 50Stoichiometry and the Mole Concept

מהי המסה המולרית של Ca(OH)₂ (Ca = 40, O = 16, H = 1 g/mol)?

  1. 57 g/mol
  2. 75 g/mol
  3. 97 g/mol
  4. 74 g/mol Correct answer.

Explanation

המסה המולרית היא סכום המסות האטומיות בכל נוסחה: 40 + 2·(16 + 1) = 40 + 34 = 74 g/mol. יש להכפיל את תת-הקבוצה OH בשתיים מפני שיש שני יוני הידרוקסיד. התוצאה 74 g/mol היא המסה של מול אחד של יחידות הנוסחה.