Atoms, Molecules and Ions Study Guide (page 3)
John Dalton's atomic theory explains the solar system-type model of an atom with electrons orbiting around a compact nucleus with protons and neutrons. Atoms make up the elements, molecules, and compounds that ultimately create chemical processes. Each substance has its own unique name and molecular formula to describe its chemical properties.
Dalton's Atomic Theory
In 1808, John Dalton published A New System of Chemical Philosophy, which proposed his hypotheses about the nature of matter. Dalton's atomic theory explained that
- all elements are made of tiny, indivisible particles called atoms (from the Greek atomos, meaning indivisible).
- atoms of one element are identical in size, mass, and chemical properties.
- atoms of different elements have different masses and chemical properties.
- compounds are made up of atoms of different elements in a ratio that is an integer (i.e., whole number) or a simple fraction.
- atoms cannot be created or destroyed. They can be combined or rearranged in a chemical reaction.
Subsequent experiments, notably those of J.J. Thomson (discoverer of the electron), E. Rutherford (who established that the atom was made of a dense, central core called a nucleus, positively charged by protons, and separated from moving electrons by empty space), and others such as A. Becquerel and Marie Curie (on the spontaneous disintegration of some nucleus with the emission of particles and radiation), were necessary, however, to complete the understanding of atoms.
The atomic weight (or mass) of an element is given by the weighted average of the isotopes' masses. Isotopes are atoms of an element that have different masses.
Dalton's atomic theory allowed scientists to understand and formulate three laws based on Dalton's hypothesis:
|1.||Law of conservation of mass: The law of conservation of mass is derived from Dalton's fifth hypothesis and states that mass cannot be created or destroyed. If the mass of the combined reactants is 20 grams, then the mass of the combined products must be 20 grams.|
|2.||Law of definite proportions: The law of definite proportions is derived from Dalton's fourth hypothesis and states that different samples of the same compound always contain the same proportion by the mass of each element.Water (H2O) always has a ratio of 2 grams of hydrogen to 16 grams of oxygen regardless of the sample size.|
|3.||Law of multiple proportions: The law of multiple proportions is also derived from Dalton's fourth hypothesis and states that if two elements combine to form multiple compounds, the ratio of the mass of one element combined with the 1 gram of the other element can always be reduced to a whole number. Hydrogen can combine with oxygen in two ways: water (H2O) and hydrogen peroxide (H2O2). The ratio of oxygen is 1:2.|
Elements are defined by their atomic number. An element's atomic number is the number of protons in the atom and is sometimes written as a subscript of the elemental symbol (i.e., 11Na). Because the mass number defines the elemental symbol (sodium always has 11 protons and carbon always has 6 protons), the atomic number is frequently omitted.
Also important is the mass number of an element. The mass number is the sum of protons and neutrons (in the nucleus) of the atom and is written as a superscript of the element's symbol (23Na).
Isotopes are atoms of the same element with the same number of protons (same atomic number) but different number mass numbers (due to a different number of neutrons). Isotopes have identical chemical properties (the same reactivity) but different physical properties (i.e., some are radioactive, while others are stable). Consider the three isotopes of hydrogen in Table 2.1.
Elements usually form +1 cations in group 1,+2 cations in group 2, and +3 cations in group 3.
Nonmetals usually form –1 anions in group 7,–2 cations in group 6, and –3 anions in group 5.
As noted earlier, the periodic table is organized in octaves (Groups 1A to 8A). The octet rule states that atoms form ions and are bond by surrounding themselves with eight (octet) outer electrons. Notable exceptions are hydrogen (two electrons; duet rule) and group 3 elements (six electrons). They tend to acquire the stability of their closest noble gases in the periodic table either by losing (metals), gaining (nonmetals), or sharing electrons in their valence shell. The valence shell contains the electrons in the outermost energy level.
An anion is a negatively charged ion formed when an atom gains one or more electrons. Most anions are nonmetallic. Their names are derived from the elemental name with an -ide suffix. For example, when chlorine (Cl) gains an electron, a chloride ion (Cl–) is formed. Because chlorine is in group 7, it only needs to gain one electron to achieve the octet structure of argon (Ar). An oxygen atom (O) acquires two electrons in its valence shell to form an oxide ion (O2–) that has the same stable electron configuration as neon (Ne).
A cation is a positively charged ion formed when an atom loses one or more electrons. Most cations are metallic and have the same name as the metallic element. For example, when lithium (Li) loses an electron, a lithium ion (Li+) is formed. Lithium is in group 1 and needs to lose one electron to acquire the noble gas electron structure of helium (He), the closest noble gas.
Ionic compounds are compounds formed by combining cations and anions. The attractive electrostatic force between a cation and an anion is called an ionic bond.
A molecular compound is formed when a nonmetal and metal combine to form a covalent bond. Covalent bonds are the type of bonds formed when two atoms share one or more pairs of electrons to achieve an octet of electrons. A polar covalent bond is formed when the atoms unequally share paired electrons.
Electronegativity is the ability of an atom (in a bond) to attract the electron density more than the other atom(s). Electronegativity increases from left to right and from the bottom to the top of the periodic table. Thus, fluorine (F) is the most electronegative element of the periodic table, with the maximum value of 4.0 in the Pauling scale of electronegativity. Metals are electropositive. See Lesson 11 for more on electronegativity.
Formulas and Nomenclature
An essential step in learning chemistry is understanding chemical formulas and how to name compounds. Compounds can be divided into four classifications:
|1.||Type I: binary ionic compounds|
|2.||Type II: binary ionic compounds with the metal possessing more than one type of cation|
|3.||Type III: binary covalent compounds|
Type I and II Binary Compounds
Type I and II binary compounds are neutral, ionic compounds that contain two parts: a cation and an anion. When a metal is the cation and a nonmetal is an anion, the following rules are used:
- The cation is always listed first and the anion second.
- For cations that possess multiple ions (see Figure 2.1), the charge on the ion must be specified by using a Roman numeral in parentheses following the cation.
Type I and II compounds are neutral and charges must balance to create a net zero charge.
Polyatomic ions are ions that contain more than one atom (see Table 2.2). These ions can replace one or both ions in Type I or II ionic compounds and have special names. However, the oxyanions (the ions containing oxygen) have a systematic naming structure. When two oxyanions of an element are present, the anion with the larger number of oxygen atoms is given the suffix -ate (i.e., sulfate, SO4 2–, and nitrate,NO3–), and the anion with the smaller number of oxygen atoms is given the suffix -ite (i.e., sulfite, SO32–, and nitrite, NO2–).When more than two oxyanions exist in a series, the prefixes hypo- (less than) and per- (more than) are used, as in chlorine and bromine oxyanions.
Type III: Binary Covalent Compounds
Type III binary compounds are neutral, covalent compounds that contain two nonmetals. Type III naming is similar to Type I and II using the following rules:
- The element listed first is named first using the full element name.
- The element listed second is named as if it were an anion.
- A prefix is used to represent the number of atoms because nonmetals can combine in many different ways (see Table 2.3). The prefix mono- is not used for the first element.
Some covalent molecules use a common name over their systematic name. Examples include H2O (water), NH3 (ammonia), CH4 (methane), and BH3 (borane).
Acids are substances that donate positive hydrogen ions (H+) when dissolved in water. Adding one or more hydrogens to an anion requires a different name.
For anions ending in -ide, use a hydro- prefix and an -ic acid ending. For example, what is the chemical name for HCl and HCN?
For oxyanions (polyatomic ions containing oxygen), use an -ic acid ending for polyatomic ions ending in -ate and -ous ending for polyatomic ions ending in -ite. For example, what is the chemical name for H2SO4 and H2SO3?
Practice problems for these concepts can be found at - Atoms, Molecules and Ions Practice Questions
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