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0:01
Hello. This is a revision video for AP Chemistry, Chapter One: Atomic Structure and Properties. As it's a revision video, we just go through the essential information that you need to know. For more detailed explanations and full video lessons, please check the links in the description below. Good luck revising. In chemistry, to connect the amount of a substance we can measure with the number of particles it contains,
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we use a unit called the mole. A mole is defined as 6.022 times 10 to the power 23 particles of something. This is known as Avogadro's number. Just like one dozen equals 12 of something, one mole equals 6.022 times 10 to the 23 of something. Now these somethings could be atoms, molecules, or formula units, depending on the substance or the nature of what you are studying. The mole is really useful because we can measure the mass of a substance in grams quite easily,
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and then link this mass to the number of particles within the substance being studied. The molar mass of a substance in grams per mole is the mass of one mole's worth of its particles. The numerical value of molar mass in grams per mole is the same as the average mass of a single particle in atomic mass units, amu. For example, carbon has an average atomic mass of 12.01 amu, so its molar mass is 12.01 g per mole.
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This means that 12.01 g of carbon contains exactly one mole's worth of atoms, or, to be specific, 6.022 times 10 to the 23 atoms. Dimensional analysis is a method for converting between different units using conversion factors. It is essential for solving problems involving mass, moles, and particle numbers. For example, moles equals mass divided by molar mass, with mass being in grams and molar mass being in grams per mole. We can also rearrange this equation to give us
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molar mass equals mass divided by moles and mass equals moles times molar mass. Also, number of particles equals moles times Avogadro's constant. This we can also rearrange to give moles equals number of particles divided by 6.022 times 10 to the 23, or Avogadro's constant. 6.022 times 10 to the 23 equals number of particles divided by moles. Exam tip here: make sure you are comfortable using and rearranging these equations.
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Try really hard to avoid using things like formula triangles and instead understand how to rearrange these equations using basic mathematical principles. It will really help you later in the course with more complicated equations and calculations. Many elements in nature exist as mixtures of isotopes. Isotopes are atoms of the same element, but with different masses. Mass spectrometry is a technique used to separate and measure the different isotopes in a sample of an element.
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The machine used in mass spectrometry, a mass spectrometer, produces a graph called a mass spectrum. The x-axis shows the relative mass of the samples being detected, and the y-axis shows the relative abundance of each isotope in the sample, usually as a percentage. Each bar in the graph represents an isotope of the element. A taller bar means the isotope is more abundant and has a higher relative abundance.
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Exam tip here. If you are given a mass spectrum with just bars and no percentages, you can treat the bar heights as relative values, add the total height of all the bars, then calculate each bar's percentage as a fraction of the total. The m/z value just refers to mass-to-charge ratio of the ions being detected. Because the charge of most ions in mass spectrometry is only one, then this simply means mass divided by one, so the m/z value actually equals
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the relative mass of the samples being detected. Don't worry if that seems a little bit confusing, because in AP Chemistry we only focus on singly charged ions. In mass spectrometry, each peak in a spectrum represents a different isotope of the same element. Average atomic mass is given on the periodic table of elements and aren't simply averages based on isotopes, they're weighted averages that reflect both the mass and natural abundance of each isotope. To calculate
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average atomic mass, we use the formula average atomic mass equals isotope one mass times its fractional abundance, plus isotope two mass times its fractional abundance and so on. For example, if a sample of chlorine contains two isotopes, 34.97 with an abundance of 75.78 and 36.97 with an abundance of 24.22, we can convert these abundances to decimals. So 75.78 equals 0.7578, and 24.22 equals 0.2422. We can multiply these decimals by each isotope's mass
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to give the relative atomic mass overall. So 34.97 times 0.7578 equals 26.5, plus 36.97 times 0.2422, giving 35.46 in total. A chemical formula tells us the types and number of atoms in a substance. Some substances are made up of individual molecules like H₂O or C₆H₁₂O₆, whilst others consist of ions or atoms in fixed proportions like NaCl or Fe₂O₃. An empirical formula gives the simplest whole number ratio of elements in a compound. For example,
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the simplest whole number ratio of hydrogen to oxygen in hydrogen peroxide is 1 to 1, giving an empirical formula of HO. The law of definite proportions tells us that every pure compound has the same elements in the same fixed mass ratio. So even if we have different samples of a compound, the mass ratios will always be the same. This sounds more confusing than it is. For example, water, H₂O, always contains 2 grams of hydrogen for every 16 g of oxygen, the ratio of 1 to 8 by mass.
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So, regardless of how much water you have in a sample, the ratio of hydrogen to oxygen by mass will always be 1 to 8. Or in other words, every 9 grams of water contains 1 gram of hydrogen and 8 grams of oxygen. We can determine empirical formulas from mass or percentage composition data. First, we have to convert each element's mass or percentage to moles, and then use those moles to determine the ratio of atoms of elements in the compound.
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For example, find the empirical formula for a compound with the composition by mass of 52.2% carbon, 13% hydrogen, and 34.8% oxygen. Well, these percentages by mass tell us that for every 100 g of this compound, 52.2 g would be carbon, 13 g would be hydrogen, and 34.8 g would be oxygen. We can convert these masses to moles using n equals m over M, mass divided by molar mass. For carbon, that would be 52.2 divided by 12.01,
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so 4.35 moles. For hydrogen, 13.0 divided by 1.008, giving us 12.9 moles. And for oxygen, 34.8 divided by 16, giving 2.18 moles. So this means in every 100 g of our compound, there's 4.35 moles of carbon, 12.9 moles of hydrogen, and 2.18 moles of oxygen. To get this into the simplest whole number ratio, we divide by the smallest number. The smallest number here is oxygen, 2.18. So 4.35 divided
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by 2.18 gives 2, 12.9 divided by 2.18 gives approximately 6, and 2.18 divided by 2.18 gives 1, giving us an empirical formula of C₂H₆O. Exam tip here. Always check that your final empirical formula makes chemical sense. It should always be whole numbers and reflect what's possible based on the elements involved. If your calculated ratio gives 1.5 or 1.33, multiply the ratios to clear the decimal.
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For example, times everything by 2 to get rid of 0.5, or times everything by 3 to get rid of 0.33. Substances in chemistry can be classified as pure substances or mixtures, depending on their composition and uniformity. A pure substance contains only one type of particle, be it atoms, molecules, or formula units. Pure substances have a fixed composition and consistent properties throughout the sample, because they contain only one kind of particle.
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A mixture contains two or more substances physically combined together. These substances may be elements, compounds, or both. However, they aren't chemically bonded together, they're only physically combined. The proportions of substances in a mixture can vary. For example, salt water can have different concentrations. The individual components of a mixture retain their original properties and can often be
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separated out using physical techniques such as filtration, distillation, and chromatography. For both pure substances and mixtures, we can analyze elemental composition by mass. This just means measuring the percentage of each element present in a sample. For a pure substance, these percentages should match its chemical formula exactly. For a mixture, the elemental composition will reflect the relative proportions
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of each substance present, and therefore can be different for different mixtures. For example, if a sample of magnesium oxide, MgO, is tested and found to contain 60.3% magnesium and 39.7% oxygen, we can check to see how pure this sample is. First, we can calculate the expected mass percentages for pure MgO. Well, the molar mass of MgO is 24.31 plus 16, the molar masses of magnesium and oxygen added together. This gives us 40.31 g per mole.
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The percentage of magnesium here would be 24.31 divided by 40.31 times 100, 60.3%, and the percentage of oxygen would be 16 divided by 40.31 times 100, 39.7%. So in a sample of pure magnesium oxide, we would expect 60.3% to be magnesium and 39.7% to be oxygen. Here these calculated values match the tested values, so therefore the sample of magnesium oxide can be described as pure. Exam tip here. Don't assume all measured compositions are accurate.
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Even small differences can indicate impurities, so be especially careful in questions where you're asked to justify purity based on given data. Atoms are the smallest indivisible units of an element, and consist of a small, dense, positively charged nucleus containing subatomic particles called protons and neutrons. This nucleus is surrounded by another type of subatomic particle called electrons, that are negatively charged.
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Because the nucleus of an atom is positively charged and the surrounding electrons are negatively charged, there is an electrostatic force of attraction between them. This attraction can be described by something called Coulomb's law, where F is proportional to q₁ times q₂ divided by r squared, with F being the force of attraction, q₁ and q₂ being the charges of the two particles, and r being the distance between them.
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If we model the nucleus as a positively charged particle, then this equation can be used to describe the attraction between the nucleus of an atom and the electrons that surround it. The smaller the r value is, so the closer an electron is to the nucleus, the stronger the force of attraction. Equally, the greater the nuclear charge, so the more positive charge a nucleus has, the stronger the attraction as well. Electrons can exist at different distances from the nucleus of an atom,
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because they are arranged in energy levels or shells. These are labelled by the principal quantum number n. n equals 1, n equals 2, n equals 3, and so on. The model of the atom becomes more accurate when we divide each main energy level into sublevels. These are labelled as s, p, d, and f, and they differ in both energy and shape. Not all sublevels are present in each energy level in an atom, and they're sometimes referred to as subshells.
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The n equals 1 energy level has only an s sublevel. The n equals 2 energy level has an s and a p sublevel. The n equals 3 energy level has an s, a p, and a d sublevel. And the n equals 4 energy level has an s, a p, a d, and an f sublevel. Each sublevel contains one or more atomic orbitals. An orbital is a region of space in an atom where there is a high probability of finding an electron. Each individual orbital can hold up to 2 electrons as a pair, each with opposite spin.
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Different sublevels contain different numbers and shapes of orbitals. An s sublevel contains only one orbital, that is, an s orbital. It is spherical in shape and is symmetrical around the nucleus. Because it only has one orbital, then an s sublevel can only hold a maximum of 2 electrons. A p sublevel contains 3 orbitals, which are called p orbitals. They are labelled as pₓ, pᵧ, and p𝓏. They're often described as having a dumbbell-like shape,
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although this isn't strictly speaking true. It's a nice visualization to use at this level. Each p orbital can hold up to 2 electrons, meaning a p sublevel can hold a maximum of 6 electrons. A d sublevel in an atom has 5 orbitals in and can hold a maximum of 10 electrons. The shapes of orbitals in the d and f sublevels are quite complicated, and we don't need to worry about them at this stage of the course.
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From Coulomb's law, we know the closer an electron is to the nucleus of an atom, the more attraction it has to it, and therefore the lower in energy it is. As a result, electrons fill subshells in order of increase in energy or distance from the nucleus. This is called the Aufbau principle. A 1s subshell is filled before 2s, a 2p before 3s, a 3p before 4s, and a 3d before 4p. Note that the s subshell in the fourth energy level is actually filled before
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the d subshell in the third energy level, because it's actually slightly lower in energy. Electron configurations show how electrons are arranged within atoms or ions. The notation uses a number for the shell, n, a letter for the subshell, so s, p, d, for example, and a superscript for the number of electrons. For example, oxygen has an atomic number of 8 and so has 8 electrons. This would give a configuration of 1s² 2s² 2p⁴.
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Shorthand notation uses the nearest noble gas to represent inner or core electrons. Core electrons are the inner filled shells of an atom. Valence electrons are in the outermost shell and determine chemical reactivity. For example, oxygen has 2 inner electrons in the first shell and 6 valence electrons in the second, its outermost shell, the 1s² being core electrons, and the 2s² 2p⁴ being valence electrons. When atoms become ions, they lose or gain electrons.
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Cations are positively charged ions that are formed when atoms lose their highest energy electrons. For example, sodium will lose its 3s¹ electron first to become 1s² 2s² 2p⁶, because the 3s¹ electron is the highest energy electron in the atom, so it takes the least amount of energy to remove. Anions are negatively charged ions that are formed when atoms gain electrons. The gained electrons go into the next available orbital within the atom.
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For example, when an atom of chlorine gains an electron to become a chloride ion, the extra electron goes into a space in a p orbital in the 3p subshell, giving a complete configuration of 1s² 2s² 2p⁶ 3s² 3p⁶. Transition metals provide an exception to the general rules and, when transition metals form ions, electrons are usually removed from the 4s subshell before the 3d. For example, Fe will go from [Ar] 4s² 3d⁶ to 3d⁶ when it forms an Fe²⁺
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ion. Exam tip here. When writing electron configurations, always double-check that you've filled subshells in the correct order and that your total number of electrons matches the element or ion's charge. For ions, always remove electrons from the outermost shell, highest principal energy level, first. Ionization energy is the energy required to remove an electron from an atom or ion. It's based upon the distance the electron is from the nucleus of the atom,
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the nuclear charge of the atom, and the level of shielding. Electrons further from the nucleus are easier to remove, as they have less electrostatic attraction to it. The stronger the nuclear charge, if the atom has more protons in its nucleus, the stronger the attraction between the electron and nucleus, the harder it is to remove the electron. The more inner or core electrons there are inside the atom, the weaker the attraction
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the outer electrons have to the nucleus, making them slightly easier to remove. This is referred to as inner electron shielding. As a general trend, ionization energy increases across a period, so left to right in the periodic table, and decreases down a group, so top to bottom. Photoelectron spectroscopy, or PES, measures the energy needed to remove electrons from different subshells in an atom. Specifically, it measures the binding energy of electrons,
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the energy required to remove an electron from a specific orbital. Electrons closer to the nucleus are more strongly attracted to it and therefore require more energy to remove, having a greater binding energy. The data from photoelectron spectroscopy is often presented as a photoelectron spectrum, a graph with the x-axis being binding energy in megajoules per mole or electron volts, and the y-axis being relative number of electrons.
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Each peak represents electrons being removed from a specific subshell. The relative heights of the peaks show how many electrons are in that subshell, and the position of the peak, so how far on the x-axis it is, shows how strongly the electrons are attracted to the nucleus of the atom. The further right on the x-axis the peak is, the less attraction it has to the nucleus, meaning the subshell is further from the nucleus of the atom.
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A PES spectrum can be directly linked to electron configurations. Each group of peaks matches the electrons in different subshells. For example, an unknown element X has a peak spectrum with the following features. By looking at the relative number of electrons from the peak height, we can see that the first 2 peaks each refer to 2 electrons being removed and the biggest peak 5 electrons. Well, the first peak must be a 1s, the second peak must be from a 2s, and the final peak
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must be from a 2p, meaning 2 electrons in the 1s, 2 electrons in the 2s, and 5 electrons in the 2p. This would give 9 electrons in total, which would correspond to the atomic number of fluorine, meaning the unknown element is fluorine. Exam tip here. In PES questions, be ready to match spectra to electron configurations or explain differences in binding energy. Focus on how close the electrons are to the nucleus and how many occupy each subshell. Also remember that
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the x-axis is binding energy, it decreases to the right. The periodic table in chemistry shows repeating patterns in chemical and physical properties of the elements, known as periodicity. These patterns are explained by electron configurations and the arrangement of electrons in energy levels and sublevels. Elements in the same group or vertical column have similar outer electron configurations, giving them similar chemical properties.
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They have the same valency, number of electrons in their outer shell. Across a period or horizontal row, the number of protons and electrons increases, affecting atomic properties in predictable ways. These trends are largely due to the number of electron shells increasing down a group, the effective nuclear charge experienced by the outermost electrons, which increases across a period, and shielding by inner electrons, which also increases down a group.
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These all link together to affect the strength of attraction between the nucleus and valence electrons in an atom, and can be explained using Coulomb's law, where the force of attraction is proportional to the charge of particle one times charge of particle two divided by the distance between them squared. As nuclear charge, so the number of protons in the nucleus increases, the attraction to outer electrons increases unless offset by shielding from inner electrons.
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For example, ionization energy, the energy required to remove an electron from an atom, increases as a general trend across a period from left to right, and decreases down a group from top to bottom. This is because across a period, there are more protons in an atom's nucleus, meaning a stronger effective nuclear charge and stronger attraction between the outer electrons and nucleus, making them harder to remove. The number of inner electron shells
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remains the same across the period, meaning the inner shielding doesn't change. However, down a group, more inner electron shells means more distance and more shielding between the outer electrons and the nucleus of the atom, making the electrons easier to remove because they experience less attraction. Notice here that the positive charge of the nucleus does also increase down the group, but
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it doesn't increase enough to offset the increased distance between the outer electron and nucleus. The atomic radius, so distance between the nucleus to the outermost electron, decreases across a period and increases down a group. This is because, for the same reasons as ionization energy, across a period, increase in nuclear charge pulls electrons in the atom closer to it. This decreases the overall
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size of the atom. Down a group, however, new energy levels are added, making the atom larger, increasing the inner electron shielding, moving the outer electrons further away from the nucleus. Ionic radius refers to the distance between the nucleus and outermost electron in an ion. Cations, positively charged ions, are smaller than the atoms that they form from because they experience less electron repulsion. Anions, negatively charged ions, are larger than
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the atoms they form from because they have more electrons in, causing more electron repulsion. As a result, the ionic radius trend across a period actually changes. For example, in period 3, the ionic radius decreases from Na⁺ to Al³⁺ as the nuclear charge increases whilst the number of electrons stays the same. The ions all have the same electron configuration. This greater attraction from the nucleus pulls the outer shell of the ion more tightly, decreasing the ionic radius.
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That's then followed by a noticeable jump in ionic radius between Al³⁺ and P³⁻ due to the switch from positively charged to negatively charged ions, then the ionic radius decreases again as the ions get less negatively charged, causing less repulsion in the outer shell. Notice that silicon and argon in period 3 don't readily form ions, and so don't have an associated ionic radius. Electron affinity, the energy change that occurs when a gaseous atom gains an electron,
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becomes more negative across a period and becomes less negative down a group. This can be explained as a greater attraction between the nucleus of an atom and an incoming electron as you go across a period, and a decrease in attraction as you go down a group. Note that there are some slight exceptions to this trend, which are usually explained as atoms have unfilled sublevels, so incoming electrons experience less attraction to the nucleus.
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Electronegativity is the power of an atom to attract electrons to itself. As a general trend, it increases across a period from left to right and decreases down a group from top to bottom. It's a relative value and doesn't have units, with the most commonly used scale being the Pauling scale, with fluorine having a Pauling value of 4.0, the most electronegative element on the table. And for group 1 metals having the lowest electronegativity values, 0.7 to 1.0.
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The electronegativity increases across a period because the atomic radius decreases and nuclear charge increases, but the shielding stays roughly the same. So the nucleus is able to better attract an electron towards the atom. Down the group, however, atomic radius and shielding increase, so it's harder for the nucleus of an atom to attract electrons towards itself. The effective nuclear charge, or Z_eff, of an atom, is the net positive charge
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felt by an electron after accounting for shielding by inner electrons. It helps explain the periodic trends outlined previously, like atomic radius and ionization energy, where Z_eff equals actual nuclear charge minus shielding from inner electrons. As effective nuclear charge increases, electrons are held more tightly, meaning higher effective nuclear charge means a smaller atomic radius and higher ionization energy, and a lower effective nuclear charge
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means a larger atomic radius and a lower ionization energy. Valence electrons are the electrons in the outermost shell of an atom. It is these valence electrons that are responsible for nearly all the chemical behavior of atoms of elements. The number of valence electrons an atom has can be predicted from its group number in the periodic table. Group 1 elements, the alkali metals, have one valence electron. Group 2 metals, the alkaline earth metals, have 2 valence electrons,
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and this increases across a period. For example, group 17 elements, the halogens, have 7 valence electrons, and group 18 elements, the noble gases, have 8 valence electrons, apart from helium, which has only 2. The most stable electron configuration for most atoms is to have 8 electrons in its outermost shell. This is referred to as the octet rule, but there are many exceptions to this, especially transition metals. Atoms of elements can often lose or gain electrons to achieve a
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full outer shell of electrons and satisfy the octet rule. This forms charged ions, with the charge of the ion based on the number of electrons lost or gained by the atom. For example, group 1 metals lose their one valence electron to form ions of a 1 plus charge. Group 2
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metals lose their 2 valence electrons to form ions of a 2 plus charge, and group 3 metals lose their 3 valence electrons to form ions of a 3 plus charge. Nonmetals usually gain electrons to form negatively charged ions, anions. With group 15 nonmetals forming ions of a 3 minus charge by gaining 3 electrons, group 16 nonmetals gaining 2 electrons to form ions of a 2 minus charge, and group 17 nonmetals gaining 1 electron to form ions of a 1 minus charge.
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Ions form in complementary pairs in reactions where one atom loses electrons, forming a cation, and another gains those electrons, becoming an anion. The resulting oppositely charged ions formed are then held together by strong electrostatic forces of attraction, forming an ionic compound. For example, when sodium reacts with chlorine, a sodium atom loses one electron and becomes positively charged, and the chlorine atom gains that electron and becomes negatively
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charged. These ions then end up attracted to each other, forming sodium chloride, NaCl. Note the elements in group 14 don't often form ions, but instead share electrons. Transition metals are in the middle block of the periodic table and can form ions with multiple charges due to their electron configurations involving d subshells. Exam tip here. Always try and relate the charge of an ion to its element's position in the periodic table. If you're unsure about an ion's charge,
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count the valence electrons an atom of the element has, and think about how many need to be lost or gained to get a stable configuration, often satisfying the octet rule.
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