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Lesson Video: Alkali Metals Chemistry • 7th Grade

In this video, we will learn how to describe the compounds and reactivities of alkali metals and trends in their physical and chemical properties.

16:40

Video Transcript

In this video, we will learn about the alkali metals. Specifically, we will learn which elements are classified as alkali metals, where they are located on the periodic table, and the property trends in this group of elements. We will have a look at the reactivity of the alkali metals and discuss why they react in the way that they do. Lastly, we will investigate the reaction of the alkali metals with halogens, oxygen, and water and have a look at the equations for these reactions.

Let’s begin by having a look where the alkali metals are located on the periodic table. The alkali metals is the name given to the group one elements on the periodic table. These metals are grouped together because they react in a similar manner. They are called the alkali metals because all form alkaline solutions when they react with water. The alkali metals do not include hydrogen at the top of the group because hydrogen is not a metal but a nonmetal and a gas under standard conditions. The alkali metals consist of Li, Na, K, Rb, Cs, and Fr. Let’s have a closer look at these elements.

Li is lithium, Na sodium, K potassium, Rb rubidium, Cs cesium, whose IUPAC spelling includes an A, and Fr francium. Sodium and potassium are relatively abundant in the Earth’s crust compared to the other alkali metals. Lithium, rubidium, and cesium are relatively rare, while francium is extremely rare, occurring only in trace amounts in the Earth’s crust. This element is also radioactive, and not much is known about its chemistry because it is so scarce. All the alkali metals show the typical properties of metals. When their surfaces are cleaned, they are shiny. They are ductile and can be pulled into thin wires, they are malleable and can be hammered into flat sheets, and they all conduct electricity and heat very well.

Now that we know where the alkali metals are located on the periodic table, what their names and symbols are, and a little bit about them, let’s discuss the trends in this group. Specifically, let’s have a look at some of the physical property trends in the group one metals. A trend is a general pattern which is noticeable in the properties or behavior in a group or period of elements. Now let’s discuss melting point density and hardness trends. We’ll start with melting point. The bar graph shows visually the melting points and degrees Celsius of the group one metals. Here they are rounded off without decimal places.

We can see that moving down the group from lithium to sodium to potassium to rubidium to cesium, the melting point decreases. What about francium? According to the general trend, we can predict that francium will probably have a lower melting point than cesium. Francium’s melting point is estimated to be about 27 degrees Celsius. What is surprising is that the melting points of the group one metals are in general much lower than the melting points of other metals on the periodic table. We now know that the melting point trend is a general decrease down the group. Let’s have a look at density.

The bar graph shows the densities of the alkali metals. We can see that the general trend for density is opposite to that for melting point. Moving down the group, so the density of the alkali metals increases except for potassium whose density is slightly lower than that of sodium. But what about francium? According to the general trend, we can predict that francium’s density is probably higher than cesium’s density. Francium’s density is estimated to be about 2.4 centimeters cubed, which is higher than cesium’s value.

What is interesting is that the alkali metals are relatively light and low in density compared to the other metals on the periodic table. Specifically, lithium, sodium, and potassium all have densities less than that of water, in other words, less than one gram per centimeter cubed. So these three metals float on water. Now, we know the general trend in density. But what about hardness? The alkali metals are all relatively low in hardness. They are soft enough to cut with a knife. So for these metals, we often talk about softness instead of hardness. But let’s stick to the term hardness so that we don’t get confused.

Hardness decreases down the group. Moving down the group, the metals get easier to cut up. Because in general these metals are soft, they are not used in high strength applications. We have looked at the trends and some of the physical properties of this group. But what about the chemical properties and, specifically, reactivity? Let’s investigate the reactivity of the alkali metals in some depth.

We have seen that melting point decreases down the group, density increases, and hardness decreases. These trends, as well as reactivity, can be understood by looking at the atom size. Atom size, more formally referred to as atomic radius, is the radius of an atom based on how it interacts with other atoms. And this distance is approximately equal to the distance between the nucleus and the valence electron, as shown by the double-headed green arrow. The atoms of the group one elements all have one valence electron, just like in this diagram.

Now, let’s have a look how reactivity is affected by atomic radius. We’ll look at the first five metals of the alkali metals. Although all these elements contain one electron in their valence or outer shell, the number of electrons in the inner shells increases down the group or, in this case, from left to right. In general, the more inner electrons, the bigger the atom or the larger the atomic radius. Let’s look closely at lithium and sodium as examples of this.

If we look carefully, we can see that the distance between sodium’s positively charged nucleus and its valence electron is slightly larger than the distance between lithium’s nucleus and its valence electron. If we do the same for the other metals of group one, we can see a definite trend in the atomic radius. Moving down the group or, in this case, from left to right from lithium to cesium, we can see that the distance between the nucleus and the valence electron increases.

When the atomic radius is very small, for example, in the case of lithium, there is a relatively strong, attractive force exerted on the valence electron by the positively charged nucleus because they are close to each other. And so relatively speaking, much energy is required to remove that valence electron in a reaction.

However, when the atomic radius is very large, for example, in the case of cesium, there is a relatively weaker attractive force exerted on the valence electron by the positively charged nucleus. We would imagine that this is because the valence electron is further away from the nucleus. But this is only part of the reason. Remember that there are many more protons in the nucleus of cesium than in the nucleus of lithium. So we would expect the attractive force on the valence electron of cesium to be very strong because of these many protons.

But remember, there are also many inner electrons in cesium compared to in lithium. And these many inner electrons actually repel the valence electron in cesium, decreasing the overall attractive force experienced by that valence electron. In other words, the attractive force felt by the valence electron in cesium is a function of the distance that electron is from the nucleus and the attractive force from the nucleus and the repulsive force from the inner electrons. So the overall weak attractive force on the valence electron results in less energy being required to remove that valence electron in a reaction.

So moving down the group from lithium to cesium or, in this case, from left to right, we can see that there is a trend in the energy required for a reaction. Less and less energy is required as we go down the group. The ease with which a reaction occurs increases, and so we say that the reactivity increases. Now, lithium is rather reactive with water, oxygen, and halogens, for example. Potassium and rubidium are very reactive. And at the bottom end of the group, the elements are extremely reactive. Now that we know about the reactivity trend in this group, let’s have a look at some specific reactions that these elements undergo.

Let’s investigate the reaction of the alkali metals with atmospheric oxygen, water, and halogens. The group one metals react rapidly with oxygen to produce various oxides. For this reason, these metals are stored under mineral oil to prevent their exposure to oxygen in the atmosphere. Lithium and oxygen, represented by M, react with oxygen to form metal oxides of the form M2O which are solids. Let’s look at a real example of this general equation using sodium.

4Na solid plus O2 gas reacts to give 2Na2O solid. These reactions are highly exothermic. Sodium, potassium, rubidium, and cesium can react with oxygen to form metal peroxides of the form M2O2 which are solids. Let’s write a specific example of this equation using sodium again. And we get 2Na solid plus O2 gas reacting to form Na2O2 solid. Sodium, potassium, rubidium, and cesium can also react with oxygen to form compounds called super oxides under certain conditions. Super oxides have the general formula, MO2. A specific example of a reaction equation where a super oxide is formed has not been shown here. However, one example of these super oxides is NaO2 which is sodium super oxide.

Now that we know the two main equations for the reaction of the alkali metals with oxygen, let’s move on to their reaction with water. The alkali metals all react easily with water and rather vigorously. The picture drawn at the beginning of this lesson of the newspaper headline and the big explosion is a true story. It is a real-life example of the vigorous and even violent nature of the reaction between an alkali metal and water. All the alkali metals react in the same way with water. They all produce a metal hydroxide solution and hydrogen gas. Let’s look at a specific example of this equation.

Using lithium, we get 2Li solid plus 2H2O liquid reacting to form 2LiOH aqueous plus H2 gas. The metal hydroxides formed in this reaction are all of the form MOH. These metal hydroxide solutions are all alkaline or basic with a pH higher than seven. When universal indicator is added to any of these metal hydroxide solutions, it turns purple, showing us that these solutions are alkaline. And this fact is the very reason why these metals are called alkali metals.

Knowing that the reactivity of the alkali metals increases down the group, it makes sense to us that lithium fizzes on water when it reacts, that sodium melts, forms a ball, fizzes rapidly, and sometimes ignites with a flame, that potassium fizzes rapidly as it reacts, ignites with sparks and sometimes even a small explosion, and rubidium and cesium violently explode when reacting with water.

Let’s look at one last type of reaction, and that is with halogens. Alkali metals easily react with halogens to form metal halides, whose general formula is M+X-. Metal halides are ionic because they are composed of ions. Let’s look at a specific example of an equation. When lithium reacts with chlorine, which is an example of a gaseous halogen, we get 2Li solid plus Cl2 gas reacting to form 2LiCl solid. LiCl or lithium chloride is a solid, as are all the metal halides. These reactions are also exothermic. Now that we have looked at some of the types of reactions the alkali metals undergo, let’s summarize everything we have learned.

In this video, we have learned about the alkali metals. We learned that they are located in group one of the periodic table and include the elements lithium, sodium, potassium, rubidium, cesium, and francium but not hydrogen. We learned that they show specific trends going down the group. In terms of physical properties, we saw that going down the group, there is a decrease in the melting point, an increase in the density, and a decrease in the hardness. In terms of chemical properties, we saw that there is an increase in the reactivity moving down the group. We learned that reactivity is a function of the atomic radius and the ease with which or the amount of energy required to remove a valence electron in a reaction.

Lastly, we looked at the reaction of the alkali metals with oxygen, water, and halogens. When alkali metals react with oxygen, they form metal oxides of the form M2O or peroxides of the form M2O2. We briefly mentioned that some alkali metals react with oxygen to form super oxides. When reacting with water, the group one metals form alkaline metal hydroxide solutions, where the metal hydroxide has general formula MOH, as well as hydrogen gas. And finally, we saw that the alkali metals react easily with halogens to form ionic metal halides of general formula M+X−.

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