You could start the exploration of atoms with the ideas of Democratus BC , who believed that all matter in the universe was made up of tiny, indivisible, solid objects. Different spheres made up the different elements. One of the key problems for students learning about atoms, is that atoms are small.
Really, really small. This makes it difficult for students to conceptualise atoms as they cannot be seen, or touched, or investigated directly.
A good starting point to introduce atoms and illustrate their small size is to ask students to break up a piece of graphite the element carbon into as many small pieces as they can. No matter how many pieces the students break the graphite into, they will never get a single carbon atom. You can challenge higher attaining students to measure the size of an individual atom using this experiment from Practical physics.
You can challenge higher attaining students to measure the size of an individual atom using this experiment from Practical physics bit. When atoms combine, molecules are formed. For a few elements, when atoms of that element combine, a molecule of that element is formed eg H 2 and O 2. When atoms of some different elements combine, a molecule of a compound can form, eg H 2 O. How to teach elements and compounds , in the 11—14 series, describes different strategies for teaching elements and compounds and the common misconceptions students may hold.
How to teach elements and compounds rsc. Particle diagrams can be used to help the students visualise the difference between an atom, a molecule of an element and a molecule of a compound. In fact even Dalton in the s proposed a series of diagrams to represent the elements and compounds known at the time. Use of colour helps to distinguish between the atom types further.
Venn diagrams help students organise their understanding of the different particle types, as described in Atoms, elements, molecules, compounds and mixtures. In fact even Dalton in the s proposed a series of diagrams to represent the elements and compounds known at the time Figure 1. Venn diagrams help students organise their understanding of the different particle types, as described in Atoms, elements, molecules, compounds and mixtures rsc.
Ionic compounds are usually formed when a metal reacts with a nonmetal or a polyatomic ion. Covalent compounds are formed when two nonmetals react with each other. Since hydrogen is a nonmetal, binary compounds containing hydrogen are also usually covalent compounds.
The charge on the cation is the same as the group number. The cation is given the same name as the neutral metal atom. These elements usually form ionic compounds; many of them can form more than one cation. The charges of the common transition metals must be memorized; Group IV and V metal cations tend to be either the group number, or the group number minus two. Many of these ions have common or trivial names formed from the stem of the element name the Latin name in some cases plus the ending -ic or -ous.
The systematic names also known as the Stock system for these ions are derived by naming the metal first, followed in parentheses by the charge written in Roman numerals. For the metals below that typically form only one charge, it is not usually necessary to specify the charge in the compound name. The charge on the anion is the group number minus eight. The anion is named by taking the element stem name and adding the ending -ide.
Polyatomic ions are ions that are composed of two or more atoms that are linked by covalent bonds, but that still have a net deficiency or surplus of electrons, resulting in an overall charge on the group.
Hydrogen can participate in either ionic or covalent bonding. When participating in covalent bonding, hydrogen only needs two electrons to have a full valence shell. As it has one electron to start with, it can only make one covalent bond. Similarly, boron has 3 electrons in its outer shell. This nonmetal typically forms 3 covalent bonds, having a maximum of 6 electrons in its outer shell.
Thus, boron can never reach the octet state. Other atoms can have expanded orbitals and accept additional covalent bonds. Two of these that are important for living systems are sulfur and phosphorus. By the octet rule, sulfur can make 2 covalent bonds and phosphorus 3 covalent bonds. Sulfur can also have expanded orbitals to accept 4 or 6 covalent bonds, and phosphorus can expand to 5 covalent bonds. In many molecules, the octet rule would not be satisfied if each pair of bonded atoms shares only two electrons.
Consider carbon dioxide CO 2. If each oxygen atom shares one electron with the carbon atom, we get the following:. This does not give either the carbon or oxygen atoms a complete octet; The carbon atom only has six electrons in its valence shell and each oxygen atom only has seven electrons in its valence shell.
Thus, none of the atoms can reach the octet state in the current configuration. As written, this would be an unstable molecular conformation. Sometimes more than one pair of electrons must be shared between two atoms for both atoms to have an octet.
In carbon dioxide, a second electron from each oxygen atom is also shared with the central carbon atom, and the carbon atom shares one more electron with each oxygen atom:. In this arrangement, the carbon atom shares four electrons two pairs with the oxygen atom on the left and four electrons with the oxygen atom on the right. There are now eight electrons around each atom.
Two pairs of electrons shared between two atoms make a double bond between the atoms, which is represented by a double dash:. Some molecules contain triple bonds, covalent bonds in which three pairs of electrons are shared by two atoms. A simple compound that has a triple bond is acetylene C 2 H 2 , whose Lewis diagram is as follows:.
A coordinate bond also called a dative covalent bond is a covalent bond a shared pair of electrons in which both electrons come from the same atom.
A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei. In the formation of a simple or ordinary covalent bond, each atom supplies one electron to the bond — but that does not have to be the case. In the case of a coordinate covalent bond, one atom supplies both of the electrons and the other atom does not supply any of the electrons.
The following reaction between ammonia and hydrochloric acid demonstrates the formation of a coordinate covalent bond between ammonia and a hydrogren ion proton. If these colorless gases are allowed to mix, a thick white smoke of solid ammonium chloride is formed. To visualize this reaction, we can use electron dot configurations to observe the electron movement during the reaction.
First recall the valence electron states for all of the atoms involved in the reaction:. On the left side of the equation to the left of the arrow are the reactants of the reaction ammonia and hydrochloric acid.
On the right side of the reaction to the right of the arrow is the product of the reaction, the ionic compound — ammonium chloride. The diagram below shows the electron and proton movement during the reaction. Once the ammonium ion has been formed it is impossible to tell any difference between the coordinate covalent and the ordinary covalent bonds, all of the hydrogens are equivalent in the molecule and the extra positive charge is distributed throughout the molecule.
Although the electrons are shown differently in the diagram, there is no difference between them in reality. In simple diagrams, a coordinate bond is shown by a curved arrow.
The arrow points from the atom donating the lone pair to the atom accepting it. Although we defined covalent bonding as electron sharing, the electrons in a covalent bond are not always shared equally by the two bonded atoms. Unless the bond connects two atoms of the same element, there will always be one atom that attracts the electrons in the bond more strongly than the other atom does, as shown in Figure 5.
A covalent bond that has an unequal sharing of electrons, as in part b of Figure 5. A covalent bond that has an equal sharing of electrons part a of Figure 5. This is a nonpolar covalent bond. This is a polar covalent bond. Any covalent bond between atoms of different elements is a polar bond, but the degree of polarity varies widely. Some bonds between different elements are only minimally polar, while others are strongly polar. Ionic bonds can be considered the ultimate in polarity, with electrons being transferred completely rather than shared.
To judge the relative polarity of a covalent bond, chemists use electronegativity , which is a relative measure of how strongly an atom attracts electrons when it forms a covalent bond. There are various numerical scales for rating electronegativity. The polarity of a covalent bond can be judged by determining the difference in the electronegativities between the two atoms making the bond.
The greater the difference in electronegativities, the greater the imbalance of electron sharing in the bond. The Pauling Scale for electronegativities has the value for fluorine atoms set at 4. Although there are no hard and fast rules, the general rule is that a difference in electronegativity less than 0. When the difference in electronegativities is large enough generally greater than about 1.
An electronegativity difference of zero, of course, indicates a nonpolar covalent bond. Examples of electronegativity difference are shown in Figure 5. The diagram above is a guide for discerning what type of bond forms between two different atoms. By taking the difference between the electronegativity values for each of the atoms involved in the bond, the bond type and polarity can be predicted.
Note that full ionic character is rarely reached, however when metals and nonmetals form bonds, they are named using the rules for ionic bonding. For example, the orientation of the two O—H bonds in a water molecule Figure 5. In short, the molecule itself is polar. The polarity of water has an enormous impact on its physical and chemical properties.
Thus, carbon dioxide molecules are nonpolar overall. The physical properties of water a and carbon dioxide b are affected by their molecular polarities. Note that the arrows in the diagram always point in the direction where the electrons are more strongly attracted. Note that the electrons shared in polar covalent bonds will be attracted to and spend more time around the atom with the higher electronegativity value.
When the polarity is equal and directly opposing, as in the case of carbon dioxide b , the overall molecule will have no overall charge. Molecular compounds have many properties that differ from ionic compounds. Some of the generalizations for this group include much lower melting and boiling points when compared with their ionic counterpoints.
For example, water H 2 O has a melting point of 4 o C and a boiling point of o C compared with NaCl that has a melting point of o C and a boiling point of 1, o C. This is because the full charges created in ionic bonds have much stronger attractive force than the comparatively weak partial charges created in covalent molecules.
Covalent compounds, on the otherhand, do not typically have such well-structured 3-dimensional shapes. Thus they tend to be more brittle and break more easily when in solid form, and many are found in liquid and gas phases.
In addition, due to their lack of charges, they tend to be poor electrical and thermal conductors. Recall that a molecular formula shows the number of atoms of each element that a molecule contains. A molecule of water contains two hydrogen atoms and one oxygen atom, so its formula is H 2 O. A molecule of octane, which is a component of gasoline, contains 8 atoms of carbon and 18 atoms of hydrogen.
The molecular formula of octane is C 8 H When writing the chemical formula the element that is the least electronegative the element that is farther left or further down within the same family group is written first while the more electronegative element is written second. You will be required to know how to name simple binary covalent compounds compounds composed of two different elements.
The elements that combine to form binary molecular compounds are both nonmetal atoms or they are a combination of a nonmetal and a metalloid. This contrasts with ionic compounds, which were formed from a metal ion and a nonmetal ion.
Therefore, binary molecular compounds are different because ionic charges cannot be used to name them or to write their formulas. Another difference is that two nonmetal atoms will frequently combine with one another in a variety of ratios. Consider the elements nitrogen and oxygen. They combine to make several compounds including:. How would someone know which one you were talking about?
Each of the three compounds has very different properties and reactivity. A system to distinguish between compounds such as these is necessary. Prefixes are used in the names of binary molecular compounds to identify the number of atoms of each element. The table below shows the prefixes up to ten. The rules for using the prefix system of nomenclature of binary compounds can be summarized as follows. Note: the a or o at the end of a prefix is usually dropped from the name when the name of the element begins with a vowel.
As an example, four oxygen atoms, is tetroxide instead of tetraoxide. Some examples of molecular compounds are listed in Table 5. Notice that the mono- prefix is not used with the nitrogen in the first compound, but is used with the oxygen in both of the first two examples. The S 2 Cl 2 emphasizes that the formulas for molecular compounds are not reduced to their lowest ratios.
The o of the mono- and the a of hepta- are dropped from the name when paired with oxide. For example:. In addition to learning about how elements join together to form bonds, it is also very important to learn about how molecules interact with other molecules around them. This type of interaction, known as an intermolecular interaction , is important for determining broader characteristics of the molecule including reactivity and function.
Intermolecular interactions between molecules are dependent on the phase that the molecule exists. A phase is a certain form of matter that includes a specific set of physical properties. That is, the atoms, the molecules, or the ions that make up the phase do so in a consistent manner throughout the phase. As mentioned in Chapter 2, science recognizes three stable phases: the solid phase , in which individual particles can be thought of as in contact and held in place defined volume and shape ; the liquid phase , in which individual particles are in contact but moving with respect to each other defined volume but, shape of the container ; and the gas phase no defined shape or volume , in which individual particles are separated from each other by relatively large distances.
Not all substances will readily exhibit all phases on the Earth. For example, carbon dioxide does not exhibit a liquid or solid phase on Earth unless the pressure is greater than about six times normal atmospheric pressure. Other substances, especially complex organic molecules, may decompose or breakdown at higher temperatures, rather than becoming a liquid or a gas.
For example, think about roasting a marshmallow. If it gets too close to the flames it will become charred and blackened, breaking down the sugar molecules inside. The sugar is not converted into the liquid or gaseous phase. Thus, water is very unique in its ability to exist on the Earth in all three phase states solid ice — liquid water — water vapor. Which phase a substance adopts depends on the pressure and the temperature it experiences. Of these two conditions, temperature variations are more obviously related to the phase of a substance.
When it is very cold, H 2 O exists in the solid form as ice. When it is warmer, the liquid phase of H 2 O is present. At even higher temperatures, H 2 O boils and becomes steam gaseous phase.
Pressure changes can also affect the presence of a particular phase as we indicated for carbon dioxide , but its effects are less obvious most of the time. We will mostly focus on the temperature effects on phases, mentioning pressure effects only when they are important. Most chemical substances follow the same pattern of phases when going from a low temperature to a high temperature: the solid phase, then the liquid phase, and then the gas phase.
However, the temperatures at which these phases are present differ for all substances and can be rather extreme. As you can see, there is extreme variability in the temperature ranges. Recall that the melting point of a substance is the temperature that separates a solid and a liquid. The boiling point of a substance is the temperature that separates a liquid and a gas.
What accounts for this variability? Why do some substances become liquids at very low temperatures, while others require very high temperatures before they become liquids?
It all depends on the strength of the intermolecular interactions between the particles of substances. Although ionic compounds are not composed of discrete molecules, we will still use the term intermolecular to include interactions between the ions in such compounds.
Substances that experience strong intermolecular interactions require higher temperatures to become liquids and, finally, gases. Substances that experience weak intermolecular interactions do not need much energy as measured by temperature to become liquids and gases and will exhibit these phases at lower temperatures. Intermolecular forces determine bulk properties such as the melting points of solids and the boiling points of liquids.
Liquids boil when the molecules have enough thermal energy to overcome the intermolecular attractive forces that hold them together, thereby forming bubbles of vapor within the liquid. Similarly, solids melt when the molecules acquire enough thermal energy to overcome the intermolecular forces that lock them into place in the solid.
Intermolecular forces are electrostatic in nature; that is, they arise from the interaction between positively and negatively charged species. Like covalent and ionic bonds, intermolecular interactions are the sum of both attractive and repulsive components.
Because electrostatic interactions fall off rapidly with increasing distance between molecules, intermolecular interactions are most important for solids and liquids, where the molecules are close together. These interactions become important for gases only at very high pressures, where they are responsible for the observed deviations from the ideal gas law at high pressures. Substances with the highest melting and boiling points have covalent network bonding.
This type of interaction is actually a covalent bond. In these substances, all the atoms in a sample are covalently bonded to the other atoms; in effect, the entire sample is essentially one large molecule. Many of these substances are solid over a large temperature range because it takes a lot of energy to disrupt all the covalent bonds at once.
One example of a substance that shows covalent network bonding is diamond Figure 5. Diamond, a form of pure carbon, has covalent network bonding. For interactions between different molecules, the strongest force between any two particles is the ionic bond , in which two ions of opposing charge are attracted to each other. Thus, ionic interactions between particles are an intermolecular interaction.
Substances that contain ionic interactions are strongly held together, so these substances typically have high melting and boiling points.
Sodium chloride Figure 5. Solid NaCl is held together by ionic intermolecular forces. Many substances that experience covalent bonding exist as discrete molecules and do not engage in covalent network bonding.
Thus, most covalently bonded molecules will also experience intermolecular forces. At this point in your study of chemistry, you should memorize the names, formulas, and charges of the most common polyatomic ions. Because you will use them repeatedly, they will soon become familiar.
Note that there is a system for naming some polyatomic ions; -ate and -ite are suffixes designating polyatomic ions containing more or fewer oxygen atoms. The nature of the attractive forces that hold atoms or ions together within a compound is the basis for classifying chemical bonding. When electrons are transferred and ions form, ionic bonds result. Ionic bonds are electrostatic forces of attraction, that is, the attractive forces experienced between objects of opposite electrical charge in this case, cations and anions.
Covalent bonds are the attractive forces between the positively charged nuclei of the bonded atoms and one or more pairs of electrons that are located between the atoms. Compounds are classified as ionic or molecular covalent on the basis of the bonds present in them. When an element composed of atoms that readily lose electrons a metal reacts with an element composed of atoms that readily gain electrons a nonmetal , a transfer of electrons usually occurs, producing ions.
The compound formed by this transfer is stabilized by the electrostatic attractions ionic bonds between the ions of opposite charge present in the compound. A compound that contains ions and is held together by ionic bonds is called an ionic compound.
The periodic table can help us recognize many of the compounds that are ionic: When a metal is combined with one or more nonmetals, the compound is usually ionic.
This guideline works well for predicting ionic compound formation for most of the compounds typically encountered in an introductory chemistry course. However, it is not always true for example, aluminum chloride, AlCl 3 , is not ionic. You can often recognize ionic compounds because of their properties. Ionic compounds are solids that typically melt at high temperatures and boil at even higher temperatures.
When molten, however, it can conduct electricity because its ions are able to move freely through the liquid Figure 3. Figure 3. Watch this video to see a mixture of salts melt and conduct electricity. Note that the video has no narration. You can access the audio description using the widget below the video. In every ionic compound, the total number of positive charges of the cations equals the total number of negative charges of the anions.
Thus, ionic compounds are electrically neutral overall, even though they contain positive and negative ions. We can use this observation to help us write the formula of an ionic compound.
The formula of an ionic compound must have a ratio of ions such that the numbers of positive and negative charges are equal. What is the formula of this compound? Figure 4. Although pure aluminum oxide is colorless, trace amounts of iron and titanium give blue sapphire its characteristic color.
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